Although the useful pH range of a buffer depends strongly on the chemical properties of the weak acid and weak base used to prepare the buffer i. The more concentrated the buffer solution, the greater its buffer capacity. If the buffer capacity is 10 times larger, then the buffer solution can absorb 10 times more strong acid or base before undergoing a significant change in pH. A buffer maintains a relatively constant pH when acid or base is added to a solution.
The addition of even tiny volumes of 0. For buffer concentrations of at least 0. Buffers function best when the pK a of the conjugate weak acid used is close to the desired working range of the buffer. This turns out to be the case when the concentrations of the conjugate acid and conjugate base are approximately equal within about a factor of For example, we know the K a for hydroflouric acid is 6.
For the weak base ammonia NH 3 , the value of K b is 1. It's always the pK a of the conjugate acid that determines the approximate pH for a buffer system, though this is dependent on the pK b of the conjugate base, obviously.
When the desired pH of a buffer solution is near the pK a of the conjugate acid being used i. HA and A minus. So the first thing we need to do, if we're gonna calculate the pH of our buffer solution, is to find the pKa, all right, and our acid is NH four plus. So let's say we already know the Ka value for NH four plus and that's 5. To find the pKa, all we have to do is take the negative log of that. So the pKa is the negative log of 5.
So let's get out the calculator and let's do that math. So the negative log of 5. Is going to give us a pKa value of 9. So pKa is equal to 9. So we're gonna plug that into our Henderson-Hasselbalch equation right here. So the pH of our buffer solution is equal to 9.
Our base is ammonia, NH three, and our concentration in our buffer solution is. We're gonna write. And that's over the concentration of our acid, that's NH four plus, and our concentration is. So this is over. So let's find the log, the log of. And so that is. So the final pH, or the pH of our buffer solution, I should say, is equal to 9. So remember this number for the pH, because we're going to compare what happens to the pH when you add some acid and when you add some base.
And so our next problem is adding base to our buffer solution. And we're gonna see what that does to the pH. So now we've added. Let's say the total volume is. So what is the resulting pH? So we're adding. So if we divide moles by liters, that will give us the concentration of sodium hydroxide. So that's our concentration of sodium hydroxide. And since sodium hydroxide is a strong base, that's also our concentration of hydroxide ions in solution.
So this is our concentration of hydroxide ions,. So we're adding a base and think about what that's going to react with in our buffer solution. So our buffer solution has NH three and NH four plus. The base is going to react with the acids. So hydroxide is going to react with NH four plus. Let's go ahead and write out the buffer reaction here. So NH four plus, ammonium is going to react with hydroxide and this is going to go to completion here.
So we're gonna make water here. And if NH four plus donates a proton, we're left with NH three, so ammonia. Alright, let's think about our concentrations. So we just calculated that we have now. For ammonium, that would be. And for ammonia it was. So let's go ahead and write 0. Compare this to the pH if the same amount of HCl is added to a liter of pure water. Therefore, the solution will contain both acetic acid and acetate ions. In this example, ignoring the x in the [C 2 H 3 O 2 — ] and [HC 2 H 3 O 2 ] terms was justified because the value is small compared to 0.
Then, we consider the equilibrium conentrations for the dissociation of acetic acid, as in Step In the presence of the acetic acid-acetate buffer system, the pH only drops from 4. A formic acid buffer is prepared with 0. The K a for formic acid is 1. What is the pH of the solution? What is the pH if 0. What would be the pH of the sodium hydroxide solution without the buffer? What would the pH have been after adding sodium hydroxide if the buffer concentrations had been 0.
The pH went up from 3. Solving for the pH of the buffer solution if 0. This shows the dramatic effect of the formic acid-formate buffer in keeping the solution acidic in spite of the added base. It also shows the importance of using high buffer component concentrations so that the buffering capacity of the solution is not exceeded.
An alkaline buffer can be made from a mixture of the base and its conjugate acid, but the formulas for determining pH take a different form. These compounds are generally weaker bases than the hydroxide ion because they have less attraction for protons.
For example, when ammonia competes with OH — for protons in an aqueous solution, it is only partially successful. Reactions with weak bases result in a relatively low pH compared to strong bases. Bases range from a pH of greater than 7 7 is neutral like pure water to 14 though some bases are greater than An alkaline buffer can be made from a mixture of a base and its conjugate acid, similar to the way in which weak acids and their conjugate bases can be used to make a buffer.
The formula for pOH is:. Weak bases exist in chemical equilibrium much in the same way as weak acids do. A base dissociation constant K b indicates the strength of the base. For example, when ammonia is put in water, the following equilibrium is set up:. Bases that have a large K b will ionize more completely, meaning they are stronger bases. As the bases get weaker, the K b values get smaller. Calculate the pH of a buffer solution consisting of 0. Privacy Policy. Skip to main content.
Acid-Base Equilibria. Search for:. Buffer Solutions Preparing a Buffer Solution with a Specific pH A buffer is a solution of weak acid and conjugate base or weak base and conjugate acid used to resist pH change with added solute.
Learning Objectives Describe the properties of a buffer solution. Key Takeaways Key Points Buffer solutions are resistant to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A-.
When some strong acid is added to a buffer, the equilibrium is shifted to the left, and the hydrogen ion concentration increases by less than expected for the amount of strong acid added. Buffer solutions are necessary in biology for keeping the correct pH for proteins to work.
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