- Acid–base titrations
- Worked example: Determining solute concentration by acid–base titration
- Titration of a strong acid with a strong base
- Titration of a strong acid with a strong base (continued)
- Titration of a weak acid with a strong base
- Titration of a weak acid with a strong base (continued)
- Titration of a weak base with a strong acid
- Titration of a weak base with a strong acid (continued)
- 2015 AP Chemistry free response 3b
- 2015 AP Chemistry free response 3c
- 2015 AP Chemistry free response 3d
- 2015 AP Chemistry free response 3e
- 2015 AP Chemistry free response 3f
- Titration curves and acid-base indicators
- Redox titrations
Titration curves and acid-base indicators
Choosing the best indicator for different titrations depending on the pH at the equivalence point. Created by Jay.
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- What about for a WA/WB titration? What would the curve look like and what pH indicators could you use?(14 votes)
- The pH curve for a WA/WB reaction would be such that an equivalence point is hard to determine, i.e. the line isn't really straight in the middle like the other acid-base combinations(5 votes)
- Could you please explain what pKa is? Just a little background would be awesome. Overall great set of videos and super helpful and clear!(4 votes)
- Ka is the acid dissociation constant, which is primarily relevant to weak acids, because they only partially dissociate to release H+.
If HA (aq) represents a weak acid, then in aqueous solution it exists in an equilibrium of,
HA (aq) ⇌ A- (aq) + H+ (aq)
This equilibrium can be expressed in terms of Ka, because HA (aq) is a weak acid.
Ka = [A- (aq)] * [H+ (aq)]
At a given temperature, Ka is a constant.
When you've calculated Ka from the formula and equilibrial concentrations of the reaction species, you can calculate pKa, because Ka and pKa are related in the same way as [H+ (aq)] and pH,
pKa = - log( Ka )
As Ka increases, pKa decreases, because they're inversely proportional.(10 votes)
- If i don' know where the equivalence point is, how will indicator help me then?(6 votes)
- You can make a good guess.
For a strong acid-strong base titration, the equivalence point is at pH 7. The pH range of phenolphthalein is about 8.3 to 10.0, but the titration curve is so steep at the equivalence point that phenolphthalein makes a good indicator.
For a strong base-weak acid titration, the equivalence point is probably near pH 9. Phenolphthalein is great for this titration.
For a strong acid-weak base titration, the equivalence point is probably near pH 5. Here you could use an indicator like methyl red (pH 4.4 to 6.2).(6 votes)
- For a SA/SB titration, I see why all three indicators can be used, but would there be any advantage in using Bromothymol Blue over either phenophtalein or methyl red? Would it be much more accurate than the other two indicators at determining when the equivalence point is reached?(2 votes)
- In theory, yes. In practice, no.
Usually, one drop of tiitrant past the equivalence point changes the pH so much that all three indicators will change colour.(6 votes)
- Does the colour goes away if past the range of the indicator? For example, if I used Bromthymol blue as an indicator, will the color disappear once the pH goes past 7.6?(3 votes)
- Most indicators I've worked with don't change color significantly after you go out of their useful range.
This appears to be also true for Bromothymol Blue, which google image searching suggests just becomes a darker shade of blue as the solution gets more basic at least as far as pH 14 -- e.g. http://kenpitts.net/bio/energy/bromthymol_small.jpg ...(3 votes)
- Why doesn't a titration of HCL against Na2CO3 go to completion when phenolphthalein is used as indicator(2 votes)
- When you titrate HCl against Na2CO3, the reaction proceeds in two steps. The first step involves the production of NaHCO3 and this step has an equivalence point of about pH 8.3, in the range covered by phenolphthalein. Hence, the phenolphthalein changes colour at this first equivalence point.
As more acid is dripped in, the NaHCO3 reacts to liberate CO2. The equivalence point for this second step is pH 3.7. However, the phenolphthalein has already changed colour so you will not detect this second equivalence point when using phenolphthalein.
The way around this is to use a different indicator, such as methyl orange, that doesn't change colour until the second equivalence point is reached, at which point the reaction has gone to completion.(4 votes)
- At2:54he says : " Because you have this really steep titration curve you could have used any of the three acid-base indicators to find the equivalence point for your titration. "
At the equivalence point pH = 7. If we used Methyl red the indicator wouldn't have changed color since its range is 4.4-6.2, no?(2 votes)
- It would change color, from red to yellow. You just probably wouldn't see any in between, i.e. it would go straight from red to yellow because of how steep the curve is. When we say that the range of the indicator is 4.4-6.2, it would still work because 6.2 is reasonably close to 7 (at least for Jay).
Hope this helps!(3 votes)
- How could we predict the equilibriant point on thethe graph？(2 votes)
- Why is the titration curve sigmoidal? Shouldn't the pH change by the same amount for each drop of titrand?(2 votes)
- Why the equivalent point is not on pH 7 sometimes? When you do a titration, you want to know the concentration of a certain solution, so the equivalent point should always at pH 7, as when it is pH 7, the acid and bases react completely.(2 votes)
- [Voiceover] Now that we've looked at titration curves in great detail, let's see how we could use an acid-base indicator to find the equivalence point for a titration. An indicator changes color in a specific pH range. For example, methyl red is an indicator that goes from red to yellow over a pH range of about 4.4 to 6.2. We could say that's a pH range of approximately 4 to 6. Bromthymol blue is an indicator that goes from yellow to blue over a pH range of about 6 to 7.6. We could say that's approximately 6 to 8. Phenolphthalein is an indicator that goes from colorless to pink or magenta over a pH range of about 8.2 to 10. We could say that's approximately 8 to 10 here. Let's look at our titration curves, and let's see which acid-base indicators we could use. All right, we'll start at the top left here. This was the titration curve for the titration that was strong acid with a strong base. At the equivalence point, the pH is equal to 7. So right here would be a pH equal to 7. This was our equivalence point, right? So we drop down here. For the example that we did, we did the math. It took 20 milliliters of our strong base to reach the equivalence point. If you're doing a titration of a strong acid with a strong base, the pH is equal to 7 at the equivalence point. Let's think about which acid-base indicator we could use for this titration. You want to choose an indicator that changes color in a range close to your equivalence point. For example, bromthymol blue, it changes color from 6 to 8. At a pH of 6, bromthymol blue is yellow. So right here, bromthymol blue would be yellow, and then it would change to blue. You might even see some green in there. You see a color change in this range. So bromthymol blue would be a good acid-base indicator to use for your titration. When the blue color persists, you've reached the endpoint of your titration, and you've matched the endpoint of your titration with your equivalence point pretty well. This is a good way to approximate the equivalence point. For a strong acid, strong base, we have a really steep titration curve from a pH of about 4, let me change colors there, from a pH of about 4 to a pH of about 10. We could have used the other two acid-base indicators too. We could have used methyl red because methyl red changes from red to yellow in a range of about 4 to 6. That gives you a good approximation of the equivalence point. Or you could have used phenolphthalein because phenolphthalein changes to magenta somewhere between 8 to 10. Because you have this really steep titration curve like this, you could have used any of the three acid-base indicators to find the equivalence point for your titration. All right, next, let's look at the titration curve for the titration of a weak acid with a strong base. That was this example over here. We started with a weak acid. In our example, we used acetic acid, and to the acetic acid we added a strong base, sodium hydroxide. The pH at the equivalence point is greater than 7. The pH was greater than 7. For our example, it was close to 9. It was right about there, just under 9 was the pH for our equivalence point. It took about 200 milliliters of our strong base to reach our equivalence point. The reason why the pH is greater than 7 at the equivalence point for this kind of titration, is when you've neutralized all of your weak acid, you're left with a conjugate base to the weak acid. The conjugate base reacts with water to increase the concentration of hydroxide ions in solution. That's why your pH is greater than 7. Go back and watch the video for this titration if you want to see the exact calculation. For the titration of a weak acid with a strong base, the pH is greater than 7. You want to choose an acid-base indicator that changes color in a range greater than 7. For our example, phenolphthalein would work really well because it changes in a range of 8 to 10. Right in this range right in here, we would see a color change from colorless to pink or magenta, and we would stop our titration at that point, and that's a good approximation of the equivalence point. You wouldn't want to use something like methyl red here because methyl red changes in a range of 4 to 6. You could see 4 to 6 would be right in here on our titration curve. So methyl red would be red, and then it would change to yellow somewhere in here. You'd get a color change somewhere in here, so you would miss the equivalence point. This would not be a good acid-base indicator to use. So you couldn't use methyl red here. Finally, let's move on to the titration curve, for the titration of a weak base with a strong acid. We're starting with a weak base. That's why our pH is in the basic range before we've added any of our acid here. In this case, the pH at the equivalence point is less than 7. The pH of the equivalence point is less than 7. That's because if you've reacted all of your weak base with your strong acids, you're protonating your weak base. In our example, we used NH3, we used ammonia. If you're protonating your weak base, then you're gonna be left with NH4+ which is acidic. At the equivalence points you have ammonium which can donote a proton to water, and so you increase the concentration of hydronium ions in solution. At the equivalence point the pH is less than 7. Once again, watch these videos if you want to see the exact calculation for that. Your pH is gonna be less than 7. In our example, we got a little bit over 5. Right about there would be the equivalence point. If you're thinking about which acid-base indicator do you want to use, you want to use one that changes color at a pH less than 7. Methyl red would be a really good one to use here. Methyl red would be yellow at a pH of about 6. Then it would change color to red here. That would be the endpoint of your titration where you see the color change for your indicator. That would be a good way to find the equivalence point. You wouldn't want to use something like phenolphthalein because phenolphthalein would change color up here. At 10 it'd pink or magenta, and then it would change color right about there, and then you would get colorless. If it changed color here, you've missed the equivalence point. Phenolphthalein would not work for this titration. Methyl red would be the best choice.