- Acid-base definitions
- Organic acid-base mechanisms
- Ka and acid strength
- Ka and pKa review
- Using a pKa table
- Using pKa values to predict the position of equilibrium
- Stabilization of a conjugate base: electronegativity
- Acid strength, anion size, and bond energy
- Stabilization of a conjugate base: resonance
- Stabilization of a conjugate base: induction
- Stabilization of a conjugate base: hybridization
- Stabilization of a conjugate base: solvation
Bronsted-Lowry and Lewis definitions of acids and bases. Created by Jay.
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- Which is the more accepted definition of an acid: Bronsted-Lowry Acid or Lewis Acid? I have been told that the Lewis definition is the more correct one.(5 votes)
- Neither definition is "more correct", they are just different. The Lewis Acid definition is a broader one, so encompasses more compounds. A subset of these compounds would also qualify as Bronsted-Lowry Acids, since this definition is much more restrictive. There is no true definition of an acid; a chemical doesn't know if it is an acid or a base or neither. It's useful for us to classify and group compounds so that we can predict the chemical reactions that they might undergo based on the reactivity of other similarly classified compounds. Since a tendency to either donate or accept electron density (or donate or accept protons) has a large effect on reactivity, we have defined two arbitrary groups as "acid" and "base" based on these properties. The different acid/base definitions draw different boundaries for which compounds are included in these groups. Each definition is equally valid, but useful in different situations.(24 votes)
- At7:41how can it be said the "HCl is accepting a pair of electrons? " On the other side of the reaction the H is with the water. Cl is now by itself and has only one more electron than it has normally, so neither does it seem fair to say that Cl is accepting a pair of electrons.(11 votes)
- in this illustration he definitely went straight to the final result, but actually to present it in a step by step it would go like this:
- First the lone pair of the O atom would attack the H atom on HCL molecule and form a bond with H -> Now H20 has turned into a e- donor right? -> H20 is Lewis Base
- Then after the H of HCL is gone, there's a single bond left which will turn into a lone pair for Cl! So here Cl just receives a lone pair which means it's a Lewis Acid.
What is missing here is probability a second arrow to illustrate how the bond between H and Cl becomes the lone pair for Cl atom. Hope this help :)(6 votes)
- At5:36, how is Boron sp2 Hybridized?(5 votes)
- Boron is only bonded to 3 things and doesn't have any lone pairs of electrons, kind of like a carbocation.(7 votes)
- How CH3SH is more acidic than H2O ?
Thank you in advance.(3 votes)
- Sulfur is a larger atom than oxygen, thus is better able to delocalize the negative charge.(6 votes)
- Why does B get a formal charge of -1?(4 votes)
- Because B (boron) has 3 valence electrons (you can easily look that up in the periodic table), and the formal charge is calculated as follows: amount of valence electrons of the atom - actual amount of electrons surrounding it (see also the previous video, Oxidation States II, between8:30and9:00). In this case, there are 4 electrons surrounding it - in each of the 4 covalent bonds (3 with the hydrogen atoms, 1 with the oxygen), the boron gets one electron and the other atom gets the other.
So 3 - 4 is -1 formal charge. The oxygen next to it normally has 6 valence electrons, but in this case there are only 5 electrons surrounding it, so 6 - 5 = +1 formal charge.(4 votes)
- what are conjugate acids n bases?(2 votes)
- The conjugate base of an acid is the acid without its acidic proton. Likewise, if you have a base, adding a proton to its basic site would give you its conjugate acid. That is, conjugate acids and bases are related through the addition or removal of a proton.(4 votes)
- what is a conjugate base?& conjugate acid?(2 votes)
- The conjugate base is the molecule or ion left after the acid has lost a proton.
The conjugate acid is the molecule or ion formed when the base has gained a proton.(4 votes)
- Could somebody please explain to me why oxygen can form more than two bonds? I mean it is usually seen having only two bonds right? What are the rules of bonding for atoms? What is the mechanism? Is there a video on here where this is explained? This is causing me some serious headache and confusion.(1 vote)
- The important thing is the number of electrons, not the number of bonds. Oxygen has 6 valence electrons, so needs two more to complete its octet. Often, oxygen achieves this by forming two covalent bonds, as you've said. However, if an oxygen that already has two covalent bonds shares one of its lone pairs of electrons in a third covalent bond (for example, with an H+), it will still have an octet. The difference will be that, since it is now sharing these 2 electrons it now only gets half a negative charge from each (that is, it is still associated with both electrons as far as it octet goes, but it has lost one full charge (1/2 charge from each electron)). Thus, overall the oxygen will have a positive formal charge.
Conversely a negatively charged oxygen atom has 7 electrons, so only needs to form one covalent bond to complete its octet.(5 votes)
- Why does the arrow go from the electron in water to the hydrogen in HCl instead of the other way around?(2 votes)
- So what really in layman's terms is the meaning of Acid and Base and also what does Bronsted-Lowry and Lewis mean??(1 vote)
- Kinda depends on the theory...
Bronsted-Lowry: acids are H+ donors, bases are H+ acceptors
Lewis: acids are lone pair acceptors, bases are lone pair donors(2 votes)
Let's talk about the acid-base definitions for Bronsted-Lowry and, also, Lewis. And we'll start with Bronsted-Lowry. So, a Bronsted-Lowry Acid is a proton donor, and a Bronsted-Lowry Base is a proton acceptor. So let's, really quickly, review what this definition means by proton. So if I look at this diagram, right here, I'm going to draw the hydrogen atom, or the most common isotope. So hydrogen has one proton in the nucleus and one electron, somewhere around our nucleus. So a negative charge, like that. And so, we would say this is hydrogen. All right? And then we put, it's one valence electron, right there, to represent the hydrogen atom, or the most common isotope. If we were to, somehow, take away this electron, we would only be left with the proton here. We'd only be left with the proton in the nucleus. And so, when we're talking about a proton, we're talking about the nucleus of a hydrogen atom, which is equal to H plus. So, no longer are we talking about the electron. So let's see how this applies to an acid-base reaction. And so we start over here with water. And then we have HCl over here on the right. Now, in this bond, between the H the Cl, one of those electrons came from the hydrogen and one of them came from the chlorine. So let me just go ahead and draw those in. So the one from the hydrogen, I'm going to put in blue here. And that's this electron from hydrogen, right here in blue. And then for chlorine, I'm going to make that electron green. So right in here, like that. And so for this acid-base reaction, a lone pair of electrons in the oxygen is going to take this proton. So just the nucleus of the hydrogen atom leaving the electron in blue behind. And that electron in blue stays behind and ends up on the chlorine. So let's go ahead and draw what we would form from that. We would have oxygen here. The oxygen had two bonds to hydrogen. And the oxygen just picked up another bond to hydrogen. And so, let me go ahead and mark those electrons. So these electrons in here, in magenta, formed a new bond with that proton. So that's this bond right here. And then we had some electrons on oxygen. Let me go ahead and make those in red. So these electrons in red on the oxygen didn't do anything. So they're still there. So they're right here. And that's going to give that oxygen a plus one, a formal charge. And so this is the hydronium ion, H3O plus. Our other product, we would also make-- we would have our chlorine, which had three lone pairs of electrons around it already. And then it picked up both of those electrons. Let me go ahead and mark them. The one in green that it had originally brought to the dot structure. And also, the one in blue, the one it took from hydrogen like that. So chlorine now has a negative charge. So it's really the chloride anion. So this would be Cl minus, like that. So let's identify our Bronsted-Lowry Acid and our Bronsted-Lowry Base for this reaction. So let's go back over here and see what happened. So the H20, the water, acted as a proton acceptor. It accepted a proton from HCl. So water would be our Bronsted-Lowry Base. And HCl donated a proton to water. So HCl would therefore be our Bronsted-Lowry Acid. So let's go ahead and identify conjugate acid-base pairs here. So if HCl is our Bronsted-Lowry Acid, I could think about its conjugate base over here would be the chloride anions. So this would be the conjugate base over here. So H2O was our Bronsted-Lowry Base, and then over here, we can find its conjugate acid, that's H3O plus. So this would be the conjugate acid, over here. So when you're looking for conjugate acid-base pairs, you're looking for one proton difference. So H2O and H3O plus are a conjugate acid-base pair. And HCl and Cl minus are a conjugate acid-base pair. And if we look at what we have in the right here, we are now saying H3O plus is an acid, and Cl minus is a base. And so, one thing you'd think about is H3O plus donating a proton to Cl minus. And so, we'll draw a little, tiny arrow going back to the left. Because the equilibrium for this reaction lies far to the right. So we're going to get a lot more of your products on the right here. But just thinking about these definitions, right, H3O plus would be donating a proton, and Cl minus would be accepting a proton. The chloride anion would be accepting a proton. But again, we know HCl is a strong acid, so we know the equilibrium lies far to the right. So that's the idea about Bronsted-Lowry. Let's look at another definition, which is actually a little bit more broad. So this is a Lewis Acid and Lewis Base. So a Lewis Acid is an electron pair acceptor. And so, an easy way to remember this is, acid acceptor. And a Lewis Base is an electron pair donor. And so, one way to remember that this Lewis Base is an electron pair donor is to, if you think about this b being lowercase. And then just flipping it around, you would get a d here. So you get d. So a base is a donor. So let's look at this reaction here. And we have this cyclic ether, over here on the left. And then we have borine over here on the right. Now, notice there's no octet of electrons around boron, right? Boron is only surrounded by six electrons here. And that makes it very reactive. Boron is SP2 hybridized, which means it has an empty p orbital. And so, let me go ahead and represent the empty p orbital like this. It's able to accept a pair of electrons. And the ether over here is going to donate a pair of electrons. And so, let's go ahead and show what happens. The oxygen here is going to donate a pair of electrons into the empty orbital. And there's going to be a bond that forms between the oxygen and the boron. So the ether over here is donating a pair of electrons. So that must be our Lewis Base. And borine, over here, is accepting a pair of electrons. So that's our Lewis Acid. Let's go ahead and draw the product for our Lewis acid-base reaction here. So we have our oxygen is now bonded to the boron. The boron is still bonded to three hydrogens, so we draw those in there like that. And let's follow some of our electrons here before we finish drawing everything in. So these electrons in magenta formed this bond between the oxygen and the boron. And then we also had some other electrons on that oxygen. Let me identify those. So these electrons right here in red are still on that oxygen. So they are right here on that oxygen. That oxygen therefore, has a plus one, a formal charge. So plus one formal charge on oxygen. And boron gets a negative one formal charge now like that. And so, that's one Lewis acid-base reaction here. Now the Lewis acid-base definition is, once again, more inclusive than Bronsted-Lowry. If we actually go up here to the previous reaction, we can actually classify these using the definition for Lewis Acid and Lewis Base. So let's look again at what's happening here. So water is donating a pair of electrons. Well, according to Lewis Base, electron pair donor. So we could say that water, we could say this is a Lewis Base. And HCl is accepting a pair of electrons. So electron pair acceptor is Lewis Acid. So we could call this a Lewis Acid. So notice, it doesn't matter what definition you use. If you use Bronsted-Lowry, this is your acid. If you use Lewis, this is your acid. Or if you use, over here for base, this is your base, according to Bronsted-Lowry. This is also a base according to Lewis. And Lewis Acid and Base also have particular importance in organic chemistry because you can talk about the term Lewis Acid as being synonymous with electrophiles. So you could say this is an electrophile. And then, you could say a Lewis Base is an electron pair donor. That's a nucleophile. And nucleophile, electrophile are extremely important concepts to understand when you're talking about organic chemistry.