Your body is amazing – it's constantly producing substances that change the pH of your bodily fluids. From lactic acid to carbon dioxide to ketones, your cells are always at work. But what happens when the pH gets too high or too low?
If it gets too high, you get alkalosis – which can cause numbness and muscle spasms. If it gets too low, you get acidosis. But don’t worry, your body has a clever system called buffer solutions that helps keep your pH levels stable.
In this article, we’ll explore both acidic and basic buffer solutions and see what happens when you add extra acids or alkalis. You'll even get to see buffer solutions in action with a real-life example. We'll also talk about the buffer solution in blood and how it works.
But that's not all! We'll teach you how to carry out buffer solution calculations so you can put your new knowledge into practice. So let's dive into the world of buffer solutions and learn how they keep your body in balance.
A buffer solution is a solution that maintains a constant pH when small amounts of acid or alkali are added to it.
Buffer solutions can be acidic or basic. They are designed to keep the concentrations of hydrogen ions and hydroxide ions roughly the same by reacting with whatever substances are added to them, be it another acid or another base.
How Acidic Buffer Solutions Work
To create an acidic buffer solution, we mix a weak acid (HA) with one of its salts in solution (MA). The salt is also the acid’s conjugate base, which is formed when an acid loses a proton.
As we learned in the ‘Weak acids and bases’ article, weak acids only partially dissociate in solution. This is crucial for maintaining a constant pH if we add an alkali. For example, when HA dissociates, it looks like this:
HA ⇌ H+ + A-
Since the acid only partially dissociates, the concentrations of H+ ions and OH- ions remain very low.
On the other hand, the soluble salt MA fully ionizes in solution:
MA ⇌ M+ + A-
This is important for keeping a constant pH if we add an acid. By mixing HA and MA, we have created an acidic buffer solution that can resist changes in pH when small amounts of acid or alkali are added.
How Acidic Buffer Solutions Resist Changes in pH
When we add an alkali to the acidic buffer solution, we increase the number of hydroxide ions (OH-). This should increase the pH, but instead, the OH- ions react with the weak acid (HA) to form water and A- ions:
HA + OH- ⇌ A- + H2O
This means that the number of hydrogen ions (H+) and hydroxide ions (OH-) remains the same, so [H+] and [OH-] are unchanged. The buffer solution has successfully resisted changes in pH.
Additionally, OH- ions can be removed by reacting with the hydrogen ions produced when a small amount of weak acid HA dissociates, forming water. This reaction drives the equilibrium towards the weak acid side, causing more HA to dissociate until all the alkali is used up.
Although these equations are reversible, the equilibrium lies far to the right, meaning barely any reverse reaction occurs. This is how the acidic buffer solution can resist changes in pH when small amounts of acid or alkali are added.
How Acidic Buffer Solutions Resist Changes in pH (Continued)
When we add an acid to the acidic buffer solution, we expect the pH to decrease as the acid donates protons (H+) to the solution. However, the soluble salt MA comes into play. Any protons react with A- ions produced when the salt ionizes to form the weak acid HA:
A- + H+ ⇌ HA
Overall, the concentrations of hydrogen ions and hydroxide ions remain largely unchanged. The buffer solution has once again resisted a change in pH.
Adding an acid to an acidic buffer solution produces more of the weak acid HA. This is a reversible reaction, and the acid could dissociate back into H+ and A-, increasing [H+]. However, weak acids only partially dissociate in solution, so the actual increase in [H+] is very small. This is why the buffer solution is able to resist changes in pH when small amounts of acid or alkali are added.
Creating an Acidic Buffer by Half-Neutralizing a Weak Acid with a Strong Base
Another way to create an acidic buffer is by half-neutralizing a weak acid with a strong base. Half-neutralization occurs when half of the acid in a solution is neutralized by a base, using up half of the protons.
This creates the same solution as above – a solution containing a weak acid and its salt. However, because the acid is half-neutralized, the buffer solution has a unique property: its pKa equals its pH.
For example, take one mole of weak acid HA and half a mole of base MOH in a 1000 cm3 solution. At half-neutralization, half of the weak acid reacts with the base to form a salt, MA, and water, H2O. The equation for the number of moles before and after the reaction for each species is shown below:
HA + MOH → 1/2 MA + 1/2 H2O
After the reaction, we have half a mole of both the weak acid HA and the salt MA. The salt MA ionizes into M+ and A-, giving us half a mole of A- ions. The concentrations of HA and A- ions are equal, so [HA] = [A-].
The equation for Ka is:
Ka = [H+][A-] / [HA]
As [HA] and [A-] are equal, they cancel out, leaving us with:
Ka = [H+]
This means that the pH of the buffer solution equals the pKa of the acid.
Creating a Basic Buffer Solution and its Reaction to Extra Acids and Bases
Similar to acidic buffer solutions, basic buffer solutions are made from a weak base and one of its salts, which is the base's conjugate acid.
A conjugate acid is formed when a base gains a proton. For example, when ammonia (NH3) is mixed with ammonium chloride (NH4Cl), ammonia is the weak base and ammonium ion (NH4+) is its conjugate acid.
Ammonia partially dissociates in solution, as shown by the following equation:
NH3 + H2O ⇌ NH4+ + OH-
Ammonium chloride is a salt and fully ionizes in solution:
NH4Cl → NH4+ + Cl-
The buffer solution is formed by the equilibrium between the weak base and its conjugate acid:
NH3 + NH4+ ⇌ NH4+ + NH3
When an acid is added to the buffer solution, it reacts with the OH- ions produced by the weak base:
H+ + OH- → H2O
However, the conjugate acid, NH4+, reacts with any additional H+ ions, producing more weak base NH3:
NH4+ + H+ → NH3 + H2O
The same principle applies when a base is added to the buffer solution. The weak base reacts with the added OH- ions, producing more conjugate acid NH4+:
NH3 + OH- → NH4+ + H2O
Overall, the buffer solution is able to resist changes in pH when small amounts of acid or base are added, as the equilibrium between the weak base and its conjugate acid maintains a relatively constant pH.
When an acid is added to the basic buffer solution, the H+ ions react with aqueous ammonia (NH3) to form ammonium ions (NH4+), as shown by the following equation:
H+ + NH3 → NH4+
Additionally, a small proportion of the ammonia in solution dissociates into ammonium ions (NH4+) and hydroxide ions (OH-), so the hydrogen ions from the acid can also react with the hydroxide ions to form water:
H+ + OH- → H2O
As a result, the overall concentrations of hydrogen ions and hydroxide ions remain relatively constant, and the buffer solution is able to resist changes in pH.
This is because the buffer solution contains both the weak base (ammonia) and its conjugate acid (ammonium ion), which can react with added acid or base to maintain a relatively constant pH. The equilibrium between the weak base and its conjugate acid allows the buffer solution to resist changes in pH when small amounts of acid or base are added.
Yes, that's correct. When an alkali is added to the basic buffer solution, it reacts with the ammonium ions (NH4+) in solution to form ammonia (NH3) and water (H2O), as shown by the following equation:
OH- + NH4+ → NH3 + H2O
Overall, the concentration of hydroxide ions (OH-) remains largely unchanged, and the buffer solution is able to resist changes in pH.
An example of a buffer solution commonly used in chemistry is the ethanoic acid/sodium ethanoate buffer system. Ethanoic acid partially dissociates in solution, as shown by the following equation:
CH3COOH ⇌ CH3COO- + H+
Sodium ethanoate fully ionizes in solution, as shown by the following equation:
CH3COO- + Na+ → CH3COO-Na+
When an acid is added to the buffer solution, it reacts with the ethanoate ions (CH3COO-) from the sodium ethanoate:
H+ + CH3COO- → CH3COOH
When a base is added to the buffer solution, it reacts with the ethanoic acid (CH3COOH) to form ethanoate ions (CH3COO-) and water (H2O), or it reacts with the hydrogen ions (H+) produced when a small proportion of ethanoic acid dissociates:
CH3COOH + OH- → CH3COO- + H2O
H+ + OH- → H2O
Overall, the pH of the solution remains roughly the same, and the buffer solution is able to resist changes in pH.
In the bloodstream, a buffer system is maintained to keep the blood at a suitable pH of around 7.4. Carbon dioxide (CO2) and water (H2O) react in the blood to form carbonic acid (H2CO3), which forms a buffer solution with hydrogen carbonate ions (HCO3-). This buffer system helps to regulate the pH of the blood and prevent catastrophic consequences such as cell death, which can occur if the pH goes below 6.8 or above 7.8.
Buffer solutions are also used in industries such as shampoo production and brewing, where precise control of pH is necessary to avoid skin irritation and optimize fermentation, respectively.
Now that we understand how buffer solutions work, we can have a go at calculating their pH.
A solution contains 0.5 mol dm-3 CH3CH2COOH and 1.0 mol dm-3 CH3CH2COO-Na+. The Ka for CH3CH2COOH equals 1.35 x 10-5. Calculate the pH of the solution.
To find the pH, we need to know [H+], the concentration of hydrogen ions in solution. First of all, let’s write out the equation linking Ka and [H+]:
We can rearrange to find [H+]: Here, CH3CH2COOH is our weak acid, HA. Only a tiny proportion of it dissociates into ions in solution, and so [HA] roughly equals the concentration given in the question: 0.5 mol dm-3. In contrast, all of the CH3CH2COO-Na+ molecules ionise into CH3CH2COO- and Na+ ions. The CH3CH2COO- ion is our A- and has a concentration of 1.0 mol dm-3, as given in the question. Substituting these values into the equation gives us the following:
Not too tricky, right? Now let’s take another example, this time calculating the pH of a buffer solution formed in the reaction between a weak acid and a strong base.
250 cm3 of 0.1 mol dm-3 ethanoic acid is mixed with 50 cm3 of 0.2 mol dm-3 sodium hydroxide. Calculate the pH of the buffer solution formed. The Ka for ethanoic acid = 1.76 x 10-5.
You’ll notice that this is a neutralisation reaction. The acid (ethanoic acid, CH3COOH) will react with the base (sodium hydroxide, NaOH) to form a salt (sodium ethanoate, CH3COO-Na+) and water. Before we work out the pH of the solution, we first need to find out how many moles we have remaining of acid and salt, as we need these values for our calculation.At the start of our reaction, we have the following: Moles of ethanoic acid = 0.250 x 0.1 = 0.025 Moles of sodium hydroxide = 0.050 x 0.2 = 0.010 Remember to change your volume into dm3.
Sodium hydroxide is the limiting reagent. There is less of it, and so it will be used up first. The acid and base will react together until all of the sodium hydroxide is used up, as shown in the table below. We can find the concentration of the important species at the end of the reaction by dividing their number of moles by the total volume:Calculating the pH of buffer solutions. StudySmarter Originals
Look at the equation for Ka: Ethanoic acid is our weak acid, HA. The salt product, sodium ethanoate, dissociates into our negative ion, A-. We’ve worked out the concentrations of these in the table above. We also know the Ka value for ethanoic acid given in the question. We can therefore fill in these values in our equation:
A flow chart showing you how to calculate the pH of buffer solutions. StudySmarter Originals
Most exam boards don’t need you to know how to calculate the pH of a basic buffer solution, but if yours does, take a look at the ‘Deep Dive’ box down below to find a handy equation.
There is a much simpler formula you can use to find the pH of an acidic buffer solution. It takes a bit of deriving to form it from the equation for Ka, including using the laws of logarithms:
Rearrange to find [H+] then take logs of both sides:
Expand using log laws:
Multiply both sides by -1. The left-hand side now looks like pH, and the right-hand side features pKa:
In other words:
Try it using the values from the first example above:
Likewise, we can derive a similar equation for the pH of a basic buffer solution. It ends up looking like this:
Remember that pOH + pH = pKw. At room temperature, pKw = 14. Therefore:
Yes, those are all great key takeaways on buffer solutions! To summarize:
What is a buffer solution?
A buffer solution is a solution that maintains a constant pH when small amounts of acid or alkali are added to it.
What is an example of a buffer solution?
An example of a buffer solution is a mixture of ethanoic acid and sodium ethanoate. This is an acidic buffer solution.
What are the components of a buffer solution?
Buffer solutions can be acidic or basic, depending on their components. Acidic buffer solutions are made by mixing a weak acid and one of its salts. A basic buffer solution is made in a similar way – by mixing a weak base and one of its salts.
Why is a buffer solution used in an EDTA titration?
A buffer solution is used in an EDTA titration in order to maintain a constant pH. This is necessary because the reactions between EDTA and metal ions depend on the pH.
How do you form a buffer solution?
You form an acidic buffer solution by mixing a weak acid and one of its salts. A basic buffer solution can be made in a similar way, by mixing a weak base and one of its salts.
