- 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
- Buffer solution pH calculations
Buffer solutions maintain a stable pH by neutralizing added acids or bases. They consist of a weak acid and its conjugate base, which exchange protons and hydroxide ions to form water. The Henderson-Hasselbalch equation expresses the pH as a function of the PKa and the ratio of the base and acid concentrations.
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- Is it necessary for the buffer solution to have equal amounts of the conjugate acid-base pair?(13 votes)
- No, but that gives the best buffering capacity.
Ideally, the concentration of one should not be more than 10 times or less than one-tenth the concentration of the other.(25 votes)
- isn't H+ and H3O+ the same thing? That is what my ap chemistry teacher told us.(5 votes)
- Yes, they are just simpler representations of the real thing.
Various structures of hydrated proton of the type H⁺(H₂O)n are known all the way up to n =10 (H₂₁O₁₀⁺).
Let's keep it simple.(10 votes)
- So is a buffer that has the equal amounts of conjugate acid-base pair a neutral solution overall?(2 votes)
- Nope. Each buffer has a characteristic pH.
This can be explained by the Henderson Hasselbach equation! if [A-] = [HA] then...
pH=pKa + log [A-]/[HA]
pH=pKa + log (1)
pH=pKa + 0
pH=pKa *not*pH = 7
So buffers don't try to maintain neutrality, they try to keep the pH of the solution near their pKa.
Refer to this page: http://www.chemguide.co.uk/physical/acidbaseeqia/buffers.html(5 votes)
- At4:07, why don't we divide the H3O by the HA as well? I thought that both the H3O and the A must be divided by the denominator...?(3 votes)
- Well lets see, he did it like that because he wanted to show you how he could use the laws of logarithms to express it as a sum of logs, but as for your question consider 3 numbers A,B and C, I think multiplication is associative, so lets say (A*B)/C = (A/C)*B = (B/C)*A maybe an example would make it easier to see, though I could be wrong about this.
(2*4)/4 =2 also (2/4)*4 = 2 and (4/4)*2 = 2 maybe that helped? :3 Goodluck Buddy(2 votes)
- Doesn't having a weak acid/strong conjugate base affect the PH as well? how does it stay neutral, or close to that?(3 votes)
- say hydronium ions are added to a buffer solution and then neutralized by the conjugate base in the buffer solution leaving more of the conjugate acid than the conjugate base. Shouldn't that make the solution more basic overall? If that's true is the change in pH negligible?(2 votes)
- What happens if a buffer was made with only half as much acid compared to base? does equilibrium shift and buffer capacity lower?(2 votes)
- So what does the buffer solution consist of? I'm confused about what a buffer actually is?(2 votes)
- A buffer solution is a solution that only changes slightly when an acid or a base is added to it. For an acid-buffer solution, it consists of a week acid and its conjugate base. For a basic-buffer solution, it consists of a week base and its conjugate acid.
The main purpose of a buffer solution is just to resist the change in pH so that the pH of the solution won't be much affected when we add an acid or base into it. The added acid or base is neutralized.(1 vote)
- how to calculate the BUFFER CAPACITY??(1 vote)
- Buffer Capacity (phi) = ( [H+] (or [OH-]) added ) / (change in pH of the buffer)
Intuitively, the buffer capacity must be the power of a buffer to resist pH change. And this intuition is supported by the fact that 'phi' tends to infinity in the limit of pH tending to zero.(1 vote)
- Then why doesn't the weak acid in the buffer solution react with its conjugate base? Is it because the weak acid ionizes very little? Also, is it necessary for the pH of a buffer solution to be 7 or can it be different?(1 vote)
- Because they are already in an equilibrium state. Unless you disturbed the equilibrium (in normal cases is changing the concentration), it won't react with one another.
It is not necessary for the pH to be 7. If you have an acidic-buffer solution, your pH would be lower than 7. If you have a basic buffer solution, your pH would be higher than 7.
Hope that helps :)(1 vote)
- [Voiceover] Buffer solutions resist changes in pH and so let's think about a solution of a weak acid and its conjugate base. So here we have HA which is our generic weak acid and so the conjugate base would be A-. A buffer solution needs to have substantial amounts of both present and that's what I'm trying to represent over here. We have a beaker that has in this case equal amounts of HA and A-. To our buffer solution, we're gonna add some strong acids so a source of protons. I'm going to draw a proton in here so H+. I could have written H3O+ because H+ and H2O give you H3O+. I'm just writing H+ to make it a little bit easier to think about. If you're increasing the concentration of H+ ions, or you're increasing the concentration of hydrionium ions in solution, you would think that would decrease the pH dramatically and that would happen if you didn't have something to react with the protons in solution. We do in our buffer solution, right? We have A-. We have our base. The base is going to pick up this proton and H+ and A- are going to form HA. We're going to make more of our weak acid here. Let's see, we started out with some weak acids. Let me go ahead and draw that in here so some HA and then we're going to make more because H+ and A- give us HA so we're making more of our weak acid here. What about A-? We're gonna use one of these up to make HA so now we have only two of these left. We're decreasing the concentration of A- but what we've done is we've effectively removed protons from solution and that's how a buffer is able to resist a change in pH if you add acid. What about if you add base? Next, let's think about our buffer solution and this time we're going to add a strong base. We're going to increase the concentration of hydroxide ions in solution. If we increase the concentration of hydroxide ions in solution, you would think that would increase the pH. It would except for the fact that we have an acid that can react with our base. We have some weak acid present here. The hydroxide ion is going to pick up a proton from HA so we have H+ here. OH- and H+ give us H2O so we're going to make water here. We're going to lose some of our weak acid. We started with three and now we have only two left here. If OH-, if hydroxide takes a proton away from HA, then we're left with A-. We're going to increase the amount of A- that we have in solutions. Let me go ahead and draw that in here. So once again we've effectively buffered against a change in pH because the hydroxide ions that we added reacted with the acid that was present. We've effectively removed hydroxide ions from solution. So that's the idea of a buffer solution. Let's see if we can derive an equation that will allow us to do some calculations. Next let's look at what we have down here. We know that we have in a buffer we have substantial amounts of HA and A- present. We have this reaction at equilibrium. For our equilibrium expression, concentration of products over concentration of reactions. Here we have concentration of products over our concentration of reactions once again leaving water out. Let's take the negative log of both sides of this. We're gonna take the negative log of Ka so that would be -log of Ka here and the negative log of all of this which would be equal to the negative log of all of that. I'm going to write it a little bit differently. I'm gonna put the concentration of hydronium ions out front here and then have concentration of A- over concentration of HA. I wrote it like this because it makes it a little bit easier to see a property of a logarithm. Let me just go ahead and write down the property that I'm referring to. If you have log of A times B, that's the same thing as log of A plus log of B. Here we have the negative log so the negative log of AB is equal to the negative log of A plus the negative log of B. I'm just going to write minus the log of B here. In this case, concentration of H3O plus that would be A so this would be A. This over here would be B. Let's think about what that would give us now. Let's get some more space and let's think about our log properties so on the left side we have negative log of Ka. On the right side, we would have negative log of A so that would be negative log of concentration of H3O+ minus the log of B here so that would be minus the log and B was this right here. Minus the log of A- over HA. We can keep going here because we know the negative log of KA is equal to the PKa so this is equal to the PKa. Over here we have the negative log of the concentration of hydronium ions that's equal to the pH so negative log of concentration of H3O+ is equal to the ph and then we have minus log of the concentration of A- over the concentration of HA. We can re-write this for our final equation so we're just gonna write this all for pH. The pH is equal to PKa + the log of A- over the concentration of HA here and this is called the Henderson-Hasselbalch equation. Right here is the Henderson-Hasselbalch equation. It's very useful when your doing buffer calculations. We'll look at examples of this in the next video.