In this video, we're going to begin our introduction to buffers. And so, first, we need to note that the pH of most living organisms is actually maintained right around a value of 7, or having a neutral pH. Changing the pH even slightly can actually be really, really harmful to living organisms. So, it's in the best interest of living organisms to make sure that the pH stays around 7 in its neutral range. In order to make sure that the pH stays around 7, living organisms use what are known as buffers. Buffers are defined as substances that are capable of resisting changes in the pH even when acids and bases are added to the solution. Typically, acids and bases would change the pH, but if buffers are present in the solution, then even when acids and bases are added to the solution, the substance is capable of resisting changes in pH, which means that the pH will not change. Once again, this is beneficial for living organisms because if the pH changes too much then that can be really harmful. These buffers are good things that cells use to make sure that the pH is going to resist changes and stay in that neutral range. Now, depending on the situation, buffers are capable of either decreasing or increasing the concentration of hydrogen ions in the solution. Living organisms are going to use buffers to help maintain the pH or help maintain homeostasis in regards to the pH values.

Now, if we take a look at our image down below over here on the left-hand side, notice that we're showing you a pH scale, which we know goes from 0 up to 14. We know that values of 0 are going to be acidic, and values of 14 are going to be basic. And then, of course, right in the middle with the value of 7, that is going to be neutral. Most living organisms require a pH of about 7 in order for them to survive. If the pH were to tip to either side, if too much acid was added, that could tip the scales towards the acidic side. That could be really harmful for a cell. But also, if too much base was added, that could tip the scale towards the base side and that could also be really harmful. It's in the cell's best interest to make sure that the pH is maintained and that the pH is able to resist changes so that it stays about the same.

We're showing you in the example a specific type of buffer called the bicarbonate buffer system, which is found in our blood and helps to maintain the pH of our blood. On the right-hand side, we're focusing on the bicarbonate buffer system, which includes these two molecules that we see here, HCO3− and HHCO3. Notice that on the left-hand side over here, we're showing you hydrogen ions, and remember that the pH is going to reflect how many hydrogen ions there are. If there are a lot of hydrogen ions then we would have an acidic solution. So this side over here would be the acidic side, but of course, if there are not many free hydrogen ions like over here, notice there's no H⁺ anywhere, that means that the pH is going to be high and it's going to be a basic solution.

It says that HCO3− is part of a buffer system because HCO3− is capable of accepting hydrogen ions if the hydrogen ion concentration gets too high. In a scenario when the hydrogen ion concentration gets too high, bicarbonate can act to lower the hydrogen ion concentration by accepting hydrogen ions. That is going to cause the pH to increase when HCO3− accepts hydrogen ions. However, if the hydrogen ion concentration were to get too low, then this HHCO3 over here is capable of donating hydrogen ions if it gets too low. Depending on the situation, whether the hydrogen ions are too high or too low, buffers, which would include HCO3− here and HHCO3, are capable of either decreasing or increasing the hydrogen ions in the solution. This is once again going to help maintain homeostasis in regards to the pH.

This concludes our introduction to buffers, and we'll be able to get some practice applying these concepts moving forward in our course. So I'll see you all in our next video.