Acids, bases, pH, and buffers

Even if you’ve never set foot in a chemistry lab, chances are you know a thing or two about acids and bases. For instance, have you drunk orange juice or cola? If so, you know some common acidic solutions. And if you’ve ever used baking soda, or even egg whites, in your cooking, then you’re familiar with some bases as well$^1$start superscript, 1, end superscript.

You may have noticed that acidic things tend to taste sour, or that some basic things, like soap or bleach, tend to be slippery. But what does it actually mean for something to be acidic or basic? To give you the short answer:

To see where this definition comes from, let’s look at the acid-base properties of water itself.

Hydrogen ions are spontaneously generated in pure water by the dissociation (ionization) of a small percentage of water molecules. This process is called the autoionization of water:

$\text{H}_2$start text, H, end text, start subscript, 2, end subscript$\text{O}$start text, O, end text $\text{(l)}$start text, left parenthesis, l, right parenthesis, end text $\rightleftharpoons$\rightleftharpoons $\text{H}^+$start text, H, end text, start superscript, plus, end superscript $\text{(aq)}$start text, left parenthesis, a, q, right parenthesis, end text + $\text{OH}^-$start text, O, H, end text, start superscript, minus, end superscript $\text{(aq)}$start text, left parenthesis, a, q, right parenthesis, end text

The letters in parentheses just mean that the water is liquid (l), and that the ions are in aqueous (water-based) solution (aq).

As shown in the equation, dissociation makes equal numbers of hydrogen (H$^+$start superscript, plus, end superscript) ions and hydroxide (OH$^-$start superscript, minus, end superscript) ions. While the hydroxide ions can float around in solution as hydroxide ions, the hydrogen ions are transferred directly to a neighboring water molecule to form hydronium ions (H$_3$start subscript, 3, end subscriptO$^+$start superscript, plus, end superscript). So, there aren't really H$^+$start superscript, plus, end superscript ions floating around freely in water. However, scientists still refer to hydrogen ions and their concentration as if they were free-floating, not in hydronium form – this is just a shorthand we use by convention.

So, how many water molecules in a pitcher of water will actually dissociate? The concentration of hydrogen ions produced by dissociation in pure water is 1 × 10$^{-7}$start superscript, minus, 7, end superscript M (moles per liter of water).

Is that a lot or a little? Although the number of hydrogen ions in a liter of pure water is large on the scale of what we usually think about (in the quadrillions), the number of total water molecules in a liter – dissociated and undissociated – is about 33,460,000,000,000,000,000,000,000$^{2,3}$start superscript, 2, comma, 3, end superscript. (Now there's something to think about with your next glass of water!) So, autoionized water molecules are a very tiny fraction of the total molecules in any volume of pure water.

Solutions are classified as acidic or basic based on their hydrogen ion concentration relative to pure water. Acidic solutions have a higher H$^+$start superscript, plus, end superscript concentration than water (greater than 1 × 10$^{-7}$start superscript, minus, 7, end superscript M), while basic (alkaline) solutions have a lower H$^+$start superscript, plus, end superscript concentration (less than 1 × 10$^{-7}$start superscript, minus, 7, end superscript M). Typically, the hydrogen ion concentration of a solution is expressed in terms of pH. pH is calculated as the negative log of a solution’s hydrogen ion concentration:

The square brackets around the H$^+$start superscript, plus, end superscript just mean that we are referring to its concentration. If you plug the hydrogen ion concentration of water (1 × 10$^{-7}$start superscript, minus, 7, end superscript M) into this equation, you’ll get a value of 7.0, also known as neutral pH. In the human body, both blood and the cytosol (watery goo) inside of cells have pH values close to neutral.

H$^+$start superscript, plus, end superscript concentration shifts away from neutral when an acid or base is added to an aqueous (water-based) solution. For our purposes, an acid is a substance that increases the concentration of hydrogen ions (H$^+$start superscript, plus, end superscript) in a solution, usually by donating one of its hydrogen atoms through dissociation. A base, in contrast, raises pH by providing hydroxide (OH$^-$start superscript, minus, end superscript) or another ion or molecule that scoops up hydrogen ions and removes them from solution. (This is a simplified definition of acids and bases that works well for thinking about biology. You may want to visit the chemistry section to see other acid-base definitions.)

The stronger the acid, the more readily it dissociates to generate H$^+$start superscript, plus, end superscript. For example, hydrochloric acid (HCl) completely dissociates into hydrogen and chloride ions when it is placed in water, so it is considered a strong acid. The acids in tomato juice or vinegar, on the other hand, do not completely dissociate in water and are considered weak acids. Similarly, strong bases like sodium hydroxide (NaOH) completely dissociate in water, releasing hydroxide ions (or other types of basic ions) that can absorb H$^+$start superscript, plus, end superscript.

The pH scale is used to rank solutions in terms of acidity or basicity (alkalinity). Since the scale is based on pH values, it is logarithmic, meaning that a change of 1 pH unit corresponds to a ten-fold change in H$^+$start superscript, plus, end superscript ion concentration. The pH scale is often said to range from 0 to 14, and most solutions do fall within this range, although it’s possible to get a pH below 0 or above 14. Anything below 7.0 is acidic, and anything above 7.0 is alkaline, or basic.

The pH inside human cells (6.8) and the pH of blood (7.4) are both very close to neutral. Extreme pH values, either above or below 7.0, are usually considered unfavorable for life. However, the environment inside your stomach is highly acidic, with a pH of 1 to 2. How does the stomach get around this problem? The answer: disposable cells! Stomach cells, particularly those that come in direct contact with stomach acid and food, are constantly dying and being replaced by new ones. In fact, the lining of the human stomach is completely replaced about every seven to ten days.

Most organisms, including humans, need to maintain pH within a fairly narrow range in order to survive. For instance, human blood needs to keep its pH right around 7.4, and avoid shifting significantly higher or lower – even if acidic or basic substances enter or leave the bloodstream.

Buffers, solutions that can resist changes in pH, are key to maintaining stable H$^+$start superscript, plus, end superscript ion concentrations in biological systems. When there are too many H$^+$start superscript, plus, end superscript ions, a buffer will absorb some of them, bringing pH back up; and when there are too few, a buffer will donate some of its own H$^+$start superscript, plus, end superscript ions to reduce the pH. Buffers typically consist of an acid-base pair, with the acid and base differing by the presence or absence of a proton (a conjugate acid-base pair).

For instance, one of the buffers that maintain the pH of human blood involves carbonic acid (H$_2$start subscript, 2, end subscriptCO$_3$start subscript, 3, end subscript) and its conjugate base, the bicarbonate ion (HCO$_3$start subscript, 3, end subscript$^-$start superscript, minus, end superscript). Carbonic acid is formed when carbon dioxide enters the bloodstream and combines with water, and it is the main form in which carbon dioxide travels in the blood between the muscles (where it’s generated) and the lungs (where it’s converted back into water and CO$_2$start subscript, 2, end subscript, which is released as a waste product).

If too many H$^+$start superscript, plus, end superscript ions build up, the equation above will be pushed to the right, and bicarbonate ions will absorb the H$^+$start superscript, plus, end superscript to form carbonic acid. Similarly, if H$^+$start superscript, plus, end superscript concentrations drop too low, the equation will be pulled the left and carbonic acid will turn into bicarbonate, donating H$^+$start superscript, plus, end superscript ions to the solution. Without this buffer system, the body’s pH would fluctuate enough to put survival in jeopardy.