- Acid/base questions
- Acid-base definitions
- Chemistry of buffers and buffers in our blood
- Ka and acid strength
- Autoionization of water
- Definition of pH
- Strong acid solutions
- Strong base solutions
- Weak acid equilibrium
- Weak base equilibrium
- Relationship between Ka and Kb
- Acid–base properties of salts
- pH of salt solutions
- Common ion effect and buffers
- Buffer solutions
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- Why does an acid base reaction happen at all? HCL and H2O already had its octets for all atoms, now Cl- is charged and H3O+ is charged. Why are they happier this way? Why would this reaction happen at all?(20 votes)
- There are several answers to your question...The simple one is that electrons like to be closer to the highest electronegative element. Take the combustion of methane, where CH4 and O2 --> C02 + H20 At first, Carbon and Hydrogen have similar electronegativites so the electrons are toward middle, same with O2-electrons in middle. However, in CO2 AND H20, the electrons are much closer to oxygen, this releases potential energy bc electrons are closer to the higher EN atom resulting in "MORE STABLE PRODUCTS". The complicated answer has to do with looking up free Gibbs Free energy(G), enthalpy (H) and entropy (S) which tell you whether any reaction will happen and how far it will proceed at a specific temperature.(17 votes)
- What does it mean for something to be a conjugate acid or base? How am I supposed to find out from looking at an equation?(5 votes)
- Acids and Bases are on the reactants side of an equation. There conjugate acid or base will be on the products side of the equation.
Acids will have a conjugate base and Bases will have a conjugate acid.
HCl-->H(+) + Cl(-)
You have an acid on the reactants side. This means you will want to find the conjugate base on the products side. The conjugate base will act as a base in the reverse equation. Cl(-) has the potential to take in a H(+); therefore it is the conjugate base of a acid.
Here is another example:
HBr+H2O-->H30(+) + Br(-)
HBr is the acid
H2O acts as the base because it is accepting protons
H3O(+) acts as the conjugate acid of H2O, it can't take anymore protons and will wanna give protons away in the reverse reaction.
Br(-) acts as a conjugate base of HBr, it will wanna recieve protons in the reverse reaction.
Just remember that conjugate acid and bases are essentially acids and bases in the reverse reaction.
Hope This Helps! :D(13 votes)
- Doesn't the giving of a proton completely change the atom? According to the BL definition, one is determined by giving or receiving a proton. If I am not mistaken the periodic table is based off how many protons an atom has. You can have different number of electrons or neutrons and still be considered the same atom. But once you give/receive a proton then you have completely different properties. Could someone please explain how the BL definition is possible? Thank you(1 vote)
- If you're donating or accepting a Hydrogen cation, you are essentially moving a proton. The proton is Hydrogen.(9 votes)
- So when one thing like water donates a pair of electron to BF3, then is water automatically attached to BF3? Is it always like that when other different elements donate pair of electrons? Do they become attached to each other?
Thanks! :D(5 votes)
- I think so!
Check this out: http://chemwiki.ucdavis.edu/Core/Physical_Chemistry/Acids_and_Bases/Acid/Lewis_Concept_of_Acids_and_Bases
"The reaction of a Lewis acid and a Lewis base will produce a coordinate covalent bond, as shown in Figure 1 above. A coordinate covalent bond is just a type of covalent bond in which one reactant gives it electron pair to another reactant. In this case the lewis base donates its electrons to the lewis acid. When they do react this way the resulting product is called an addition compound, or more commonly an adduct."(4 votes)
- At5:28BF3 is labeled a Lewis acid. Does this mean that it's addition to water would result in a solution of lower pH in the same way as the addition of HCl to water lowers the pH of the resulting solution? Or does a lewis acid not have to affect the pH of the solution it's added to?(2 votes)
- Lewis acids are rather a different definition of an acid. They need not (but may) have an effect on pH. In fact, many Lewis acids and bases don't even have hydrogen present at all.
Unlike the other definitions of an acid or base, the Lewis definition usually requires that you specify both the Lewis acid and the Lewis base.
But, to directly answer your question BF₃ reacts with water to form boric acid and tetrafluoroboric acid. Tetrafluoroboric acid is a strong Brønsted-Lowry acid and will, therefore, lower the pH. However, since the BF₃ would react with the water, it would no longer exist.(4 votes)
- whats octet :O this isn't meant for me is it :s(0 votes)
- It's referring to the octet of electrons in the outermost shells of an atom, a state which allows the atom to emulate the stability of a noble gas. To achieve this stable state, an atom will tend to gain or lose (share) electrons.(4 votes)
- At7:25, is the H from the HCl an acid and a base? Because it's accepting the electrons from the water but also giving away its electron to the chlorine atom to do so?(1 vote)
- It's an acid. A Hydrogen bonded to any Group 7 element will be an acid.(2 votes)
- At1:35, I understand that oxygen has 6 valence electrons and that 2 is used to bond to 2 atoms of hydrogen, leaving us with 2 lone pairs of electron.
Now, why is it that despite having 2 lone pairs, H2O doesnt have a charge, but when oxygen is bonded to a 3rd hydrogen it gains a positive charge. I'm a little confused. Thanks for your help(1 vote)
- So, atoms will have a charge if the number of valence electrons is more or less than the normal number of valence electrons. Like you said, oxygen normally has 6 valence electrons. Of these 6 electrons, two will pair up and another two will pair up leaving behind one lonely electron and another lonely electron. These two lonely electrons will each form a covalent bond with a hydrogen atom to complete their octet. Covalent bonds are made of 2 electrons, each coming from one of the two atoms. In other words, the oxygen will put forth an electron, and the hydrogen puts forth its electron, and together they form a nice stable bond. They share. NO ONE STEALS ANY ONE'S ELECTRON meaning the oxygen still has 6 valence electrons, thus it has no charge.
Now, when you form H3O+, you aren't adding a hydrogen atom (which has an electron and a proton). The oxygen and this third hydrogen aren't sharing electrons since this hydrogen HAS NO electrons. You're adding a hydrogen ion which is just a proton. Thus, you're just bringing a positive charge to the molecule. This hydrogen isn't bringing an electron to share with the oxygen the way the two other hydrogens did. You know how a single bond is made of two electrons? Well, the electrons in the bond between this hydrogen and the oxygen both come from the oxygen! In other words, the oxygen is putting its share of elctrons into the bond plus one more so it's like it's losing an electron to the hydrogen. Losing a valence electron means the molecule just gained a positive charge :)
In this picture, look at how the oxygen has its 6 electrons to itself. It has no charge. https://upload.wikimedia.org/wikipedia/commons/7/78/Acqua_Lewis.png
In this picture, the O-H bond on the far right is made of two of oxygen's electrons, not hydrogen's. The hydrogen is taking advantage of the oxygen and stealing an electron. http://antranik.org/wp-content/uploads/2011/11/hydronium-h3o-dot-structure.gif(2 votes)
- To clarify, why are Acid-Base Definitions divided into Bronsted-Lowry and Lewis definitions? Will the proton donors always be electron pair acceptors? Will proton acceptors always be electron pair donors?(1 vote)
- Good question. It is said that every Bronsted-Lowry acid/base can be classified as a Lewis acid/base. But not every Lewis acid/base will be a Bronsted-Lowry acid/base. These definitions can get confusing. Think of the specificity of acid/base definitions starting with increasing specificity: Lewis (electrons) < Bronsted-Lowry (protons) < Arrhenius (not covered, but dealing with dissociation of compound to form either H+ or -OH excess)(1 vote)
- So BF3 was originally sp2 hybridized with a free p orbital (we know this from the stearic number being 3). Once it acts as a lewis acid and forms a coordinate covalent bond with oxygen, it makes sense to me that it now has an sp3 hybridization state (because stearic number is now equal to 4). Would that be correct? Thanks :D(1 vote)
- Let's look at two definitions for acids and bases, Bronsted-Lowry and Lewis, and we'll start with Bronsted-Lowry. A Bronsted-Lowry acid is a proton donor, and a Bronsted-Lowry base is a proton acceptor. Let's really quickly review what a proton refers to. For a neutral hydrogen atom, the most common isotope has one proton in the nucleus. Here's my nucleus, here's my one proton, and one electron outside of the nucleus, so here's my electron. If we take away the electron, we're left with just that proton. We're left with the nucleus of a hydrogen atom, so we could also say this is equal to H+. When you're talking about a proton donor, that's something that's donating an H+, and a proton acceptor is accepting that H+. Let's go down here to the dot structure for HCl, and let's focus in on that covalent bond. One of those electrons came from the chlorine, so let me go ahead and draw in that electron here. One of the electrons in the bond came from the chlorine and one of the electrons in the bond came from the hydrogen, so in magenta right here. That's talking about this electron right here. HCl is going to donate a proton to water, so let's go ahead and show what happens. A lone pair of electrons on the oxygen is going to pick up this proton right here, and the electron in magenta is left behind, so these two electrons come off onto the chlorine. Let's go ahead and draw the product, so we would have, we had an oxygen bonded to two hydrogens, but the oxygen just picked up a proton, so now it's bonded to three, which gives the oxygen a +1 formal charge. Let's show those electrons. These two electrons in here in red are going to pick up this proton, and those two electrons in red are going to form this bond right here. This is H3O+ or hydronium, the hydronium ion. The chlorine becomes the chloride anions. Let's go ahead and draw that in. We had three lone pairs of electrons. We got one more lone pair, so the electron in green is on the chlorine and so is the electron in magenta. The chlorine picked up a negative charge. It becomes an anion, so we get the chloride anion here. The HCl donated a proton, so it's a proton donor. It's a Bronsted-Lowry acid. Let me go ahead and write that here. Bronsted-Lowry acid, and water accepted the proton. It accepted an H+. Water is the Bronsted-Lowry base. Let's think about the possibility of the reverse reaction. If you think about what must happen, the chloride anion must function as a base and pick up a proton from the hydronium ion here. If you think about the reverse reaction, I'll put a really tiny arrow going back in reverse, because the equilibrium lies far to the right. If you think about it in reverse, the chloride anion would function as a base. This is actually the conjugate base to HCl, so let me go ahead and write that here. Conjugate base to HCl. We have a conjugate acid-base pair here, so let me go ahead and write that over here. We have HCl and Cl-. Think about the difference between these two. There's one H+ difference between that conjugate acid-base pair. Over here, water functions as a base. Once it picks up a proton over here, it could function as an acid. This would be the conjugate acid, so let me go ahead and write that. This is the conjugate acid to our base over here, to water. We have another conjugate acid-base pair, so let me go ahead and write that over here. We have H2O, H2O and H3O+ is another conjugate acid-base pair, and once again, think about the difference between these two, between H2O and H3O+. There's one proton difference. One H+ difference. When you're dealing with the Bronsted-Lowry definitions for acids and bases, think about one proton. Let's move onto the Lewis definitions, and let's get some room down here. The Lewis definition. A Lewis acid is an electron pair acceptor, and a good way to remember that is we have an A here for acid, and then electron pair acceptor, so an A right here. A Lewis base is an electron pair donor, and a good way to remember this is if you have Lewis base, so a lower-case B for base. Over here, electron pair donor. If you take this B here and just flip it, then you would get a D. That's a nice way to remember that a Lewis base is an electron pair donor. Let's look at the dot structure here for boron trifluoride. Notice that boron does not have an octet. If we count the number of electrons, here's two, four, six electrons around it. It's not an octet. This boron right here, let me go ahead and draw this in. This boron is sp2 hybridized. If it's sp2 hybridized, it has an empty P orbital. We have an empty orbital, which I'm going to represent like this, and that empty orbital is capable of accepting a pair of electrons. Boron trifluoride is going to function as a Lewis acid here. Water has a lone pair of electrons that it can donate, so water's going to function as a Lewis base. Lone pair of electrons on the oxygen. We're going to donate that lone pair into the empty orbital here. We're going to form a bond between the oxygen and the boron. Let's go ahead and do that. The oxygen was bonded to two hydrogens already. We're going to form a bond between the oxygen and the boron, and this oxygen still has one lone pair of electrons left, which is going to give this oxygen here a +1 formal charge. The boron is bonded to three fluorines, and I'm just not going to draw in the lone pairs of electrons around fluorine to save some time here. This gives the boron a -1 formal charge. Let's show those electrons. Let me go ahead and highlight this electron pair in red here. This pair of electrons gets donated to boron to form this bond right here. Notice there's no H+ changing here, and so that's why the Bronsted-Lowry definition doesn't apply to this acid-base reaction. We have to use the Lewis definition. The Lewis definition is a little bit more broad, actually, and we can go back up here to the previous reaction, and we can look at the definition for Lewis acid and base for this one. A Lewis base is an electron pair donor, and notice that's what water is doing here as well. It's donating a pair of electrons. Not only can we say this is a Bronsted-Lowry base, we could also say this is a Lewis base. What's accepting that pair of electrons? Well, it's the protons. Let me go ahead and draw this out. HCl, you could think about HCl as being H+ and Cl-. That lone pair of electrons, this pair of electrons is being accepted by the proton, so you could say that this proton here is a Lewis acid. Let me go ahead and use a different color for that. We could say that this proton here is a Lewis acid. It's accepting a pair of electrons. You could say that HCl is a source of a Lewis acid, which is H+. These are very important definitions to understand.