Name: buffer ph and the common ion effect
Uploaded: 2020-03-26T07:22:25.000Z
Duration: 4 min 35 s
Channel: JoVE

Buffers are solutions that work towards maintaining a constant pH in a system regardless of the addition of strong acids or bases. In the absence of buffers, the addition of a strong acid or a base to a solution changes the pH significantly. Buffers are crucial in everyday life. For example, blood acts as a buffer and is able to maintain a pH between 7.35 and 7.45 since a pH above 7.8 or below 6.8 can cause death.
So, how does a buffer work? A buffer is essentially a weak acid or base and its conjugate base or acid in equilibrium with each other. Take, for example, a buffer made of acetic acid and sodium acetate. The buffer solution contains a weak acid and its conjugate salt, which dissociates to form the conjugate base acetate. The solution is acidic because the Ka of acetic acid is higher than the Kb of the conjugate base.
Remember that the dissociation constant, Ka or Kb, defines the strength of the acid or base, respectively, as it is the equilibrium constant for the dissociation of the compound in water. The dissociation constant can also be represented as pKa, or the negative log of Ka, where the smaller the pKa, the stronger the acid. The same follows for pKb, where the smaller the pKb, the stronger the base.
Returning to the acetate buffer, if we add a strong base, like sodium hydroxide, the hydroxide ions react with the hydronium ions present in the solution. This causes more acetic acid to react with water in order to return to equilibrium, thus forming more hydronium ions. If we add a strong acid, like hydrochloric acid, more hydronium ions are formed in solution. Those ions react with the acetate ions to form more acetic acid. In both cases, there is very little change in the solution pH.
Keep in mind that each buffer has a specific pH range where it can buffer a solution. Here we show a titration curve for our acetic acid buffer. The buffering range is highlighted. This region is determined by the Henderson-Hasselbalch equation, which tells us that the pH of the buffer is determined by the Ka of the weak acid and the ratio of the conjugate base to the weak acid.
Far away from this pH, the weak acid or base of the buffer is depleted, and it cannot buffer pH. So, when selecting a buffer for your application, choose one whose pKa is close to the pH desired.
Finally, all buffers have a limit to their buffering capacity, meaning the amount of acid or a base that can be added to the buffer solution before the pH changes significantly. For example, if we add too much acid to our acetic acid and sodium acetate buffer solution, we will protonate all of the acetate ions and accumulate a lot of hydronium ions, thereby lowering the pH.
Similarly, if we add too much base, we will deprotonate all of the acetic acid and accumulate a lot of hydroxide ions, thereby increasing the pH. Thus, the buffering capacity is influenced by the concentration of the weak acid and conjugate salt. So, a 1 molar acetate buffer has a higher buffering capacity than a 0.1 molar acetate buffer.
In this lab, you&#39;ll prepare and examine buffers with a wide variety of pH ranges. You&#39;ll then use your buffers to determine the pKa of a pH indicator, called neutral red, and explore the change in neutral red&#39;s pKa when riboflavin binding protein is introduced.

buffers

<h4>Buffers</h4>
When an Arrhenius acid (HA) is added to water, it dissociates into its conjugate base (A-) and a hydrogen cation (H+).
HA + H2O &#8594; H+(aq) + A-(aq)
The amount of hydrogen ions present in the solution determines the acidity of the solution, where more hydrogen ions indicate a lower or more acidic pH. Similarly, when a strong Arrhenius base (BOH) is added to water, it dissociates into its conjugate acid (B+) and a hydroxide ion (OH-).
BOH + H2O &#8594; B+(aq) + OH-(aq)
Generally, the addition of strong acids or bases to a solution changes the pH dramatically because the acid or base reacts with the water molecules in solution, increasing the concentration of either hydrogen ions or hydroxide ions. However, this change in pH can be mitigated using a buffer. Buffers are solutions that work towards maintaining a constant pH in a system, regardless of the addition of strong acids or bases.
Most often, the components of a buffer are the conjugate acid-base pair of a weak acid or weak base. For this reason, strong acids or bases that dissociate completely in water make very poor buffers and weak acids or bases that partially dissociate make better buffers. When the buffer is present, the strong acid or base does not react with the water molecules present in solution and instead reacts with the weak acid/conjugate base. This results in little to no change in the pH of the solution.
<h4>The Common Ion Effect</h4>
A buffer works through a phenomenon called the common ion effect. The common ion effect occurs when a given ion is added to a mixture at equilibrium that already contains the given ion. When this happens, the equilibrium shifts away from forming more of that ion.
For example, acetic acid (CH3COOH) dissociates slightly in water, forming the acetate ion (CH3COO-) and hydrogen ion.
CH3COOH(aq) H2O &#8660; H+(aq) + CH3COO-(aq)
If more of the acetate ion is added from soluble sodium acetate, the equilibrium position shifts to the left to form more non-dissociated acetic acid, and the concentration of hydrogen ions decreases. Here, the common ion &#8212; acetate &#8212; suppresses the dissociation of acetic acid.
A buffer must contain high concentrations of both the acidic (HA) and basic (A-) components to buffer a solution. If the amount of hydrogen or hydroxide ions added to the buffer is small, they cause a small amount of one buffer component to convert into the other. As long as the concentration of added ions is much smaller than the concentrations of HA and A- already present in the buffer, then the added ions will have little effect on pH since they are consumed by one of the buffer components. When the concentration of hydrogen or hydroxide exceeds the concentrations of the acid and its conjugate base, the buffering effect is lost, and the pH will change.
<h4>Henderson-Hasselbalch Equation</h4>
The dissociation constant, Ka, of a weak acid is calculated using the concentrations of the non-dissociated acid HA, and the concentrations of the hydrogen ions and the conjugate base, A-.
<img src="/CDNSource/lm/labs/45/45_Concepts_4.jpg" alt="Acid dissociation constant equation Ka=[H+][A-]/[HA], chemistry formula." data-altupdatedat="1754904545000" data-alt="Acid dissociation constant equation, Ka=[H+][A-]/[HA], chemistry formula." />
Higher Ka values represent stronger acids, while smaller Ka values represent weaker acids. To determine the concentration of hydrogen ions, the equation is rearranged. In this form, it is clear that the ratio of acid species to base species is important in determining the concentration of hydrogen ions, and by extension, pH.
<img src="/CDNSource/lm/labs/45/45_Concepts_5.jpg" alt="Acid dissociation equation, [H+] = Ka &#215; [HA]/[A-], chemical equilibrium concept." data-altupdatedat="1754904545000" data-alt="Acid dissociation equilibrium formula, equation for calculating hydronium ion concentration." />
Taking the negative common logarithm of both sides results in the Henderson-Hasselbalch equation.
<img src="/CDNSource/lm/labs/45/45_Concepts_6.jpg" alt="Henderson-Hasselbalch equation; pH, pKa, buffer solution, acid-base equilibrium, chemical formula." data-altupdatedat="1754904545000" data-alt="Henderson-Hasselbalch equation for calculating pH, chemistry formula illustration." />
The Henderson-Hasselbalch equation enables the calculation of buffer pH directly, without having to calculate the concentration of hydrogen ions first.
For example, it can be used to determine the pH of a 1 L buffer after adding 0.02 moles of a strong base. The strong base dissociates completely, so the concentration of hydroxyl ions added is 0.02 M. This will decrease the concentration of the acid by 0.02. Assuming that the original concentration of the acid (HA) and base (A-) components are each 0.5 M, the new concentration of base increases by 0.02 M to 0.52 M, while the concentration of acid decreases by 0.02 M to 0.48 M. Knowing the pKa of the acid component of the buffer, we can substitute these new component concentrations into the Henderson-Hasselbalch equation to determine the pH.
This is useful in determining buffer capacity, or how much strong acid or strong base can be added to a buffer without significantly affecting the pH.
<h4>Buffer Capacity</h4>
Buffer capacity is the measure of a buffer&#8217;s ability to resist pH change. This ability depends on the concentration of the buffer components, meaning the acid and its conjugate base. A higher buffer concentration has a greater buffer capacity. This means that a greater amount of hydrogen ions, or a stronger acid, would have to be added to disrupt the equilibrium and change the pH of the buffer.
Buffer capacity is also affected by the relative concentrations of the buffer components. The buffer is more effective when the concentrations of the components are similar. If the buffer component ratio is similar, then the component concentration ratio doesn&#8217;t change significantly when acid or base is added; large amounts of acid or base must be added to offset the ratio and change the pH.
The pH of the buffer differs from its buffer capacity. The pH range is the range at which the buffer is effective. Typically, buffers have a usable range within 1 pH unit of the pKa of the acid component of the buffer.
<h4>References</h4>
<ol>
	<li style="font-size:17px">Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chemistry and Chemical Reactivity. Belmont, CA: Brooks/Cole, Cengage Learning.</li>
	<li style="font-size:17px">Silberberg, M.S. (2009). Chemistry: The Molecular Nature of Matter and Change. Boston, MA: McGraw Hill.</li>
</ol>

Buffers

Buffers
When an Arrhenius acid (HA) is added to water, it dissociates into its conjugate base (A-) and a hydrogen cation (H+).
HA + H2O &#8594; H+(aq) + ...

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buffer ph and the common ion effect

Buffer solutions help keep pH steady when acids or bases are added. In water, an Arrhenius acid (HA) dissociates into hydrogen ions, H+, and a conjugate base, A-. More hydrogen ions make a solution more acidic and lower its pH. An Arrhenius base (BOH) dissociates into a conjugate acid, B+, and hydroxide ions, OH-. Strong acids and strong bases usually change pH a lot because they react completely in water and raise the amount of H+ or OH-.

Buffers reduce those pH changes. A buffer usually contains a weak acid and its conjugate base, or a weak base and its conjugate acid. Weak acids and weak bases are better buffers because they dissociate only partially in water. When a buffer is present, added strong acid or base reacts with the buffer components instead of causing a large shift in pH.

The way buffers work is called the common ion effect. A common ion is an ion already present in an equilibrium mixture. When more of that ion is added, the equilibrium shifts away from making more of it. A good example is acetic acid, CH3COOH, which dissociates slightly in water to form H+ and acetate, CH3COO-. If sodium acetate adds more acetate ion, the equilibrium shifts left and less hydrogen ion remains in solution.

A buffer needs high amounts of both the acidic form, HA, and the basic form, A-. Small amounts of added H+ or OH- are used up by one buffer component and converted into the other. If the added acid or base is much smaller than the amounts already in the buffer, the pH changes very little. Once the concentration of added H+ or OH- becomes larger than the buffer components, the buffer effect is lost and the pH shifts.

The Henderson-Hasselbalch equation helps calculate buffer pH. It comes from the acid dissociation constant, Ka, which uses the concentrations of HA, H+, and A-. A larger Ka means a stronger acid, while a smaller Ka means a weaker acid. After rearranging and taking the negative log, the equation shows that the ratio of acid to base controls pH, so buffer pH can be found without first calculating hydrogen ion concentration.

For example, a 1 L buffer with 0.5 M HA and 0.5 M A- can change after 0.02 moles of strong base are added. The strong base dissociates completely, so 0.02 M OH- is present. That amount lowers the acid concentration to 0.48 M and raises the base concentration to 0.52 M. The new values can be placed into the Henderson-Hasselbalch equation with the acid pKa to find the pH. This also helps estimate buffer capacity, or how much strong acid or strong base a buffer can absorb without a large pH change.

Buffer capacity depends on concentration. A more concentrated buffer has a greater ability to resist pH change. It also depends on how similar the acid and base concentrations are. Buffers work best when those concentrations are close, because the ratio changes only a little when acid or base is added. Buffer pH is not the same as buffer capacity. The effective pH range is usually within 1 pH unit of the acid component's pKa.

Buffer pH and the Common Ion Effect

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Buffer pH and the Common Ion Effect

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