One topic in chemistry that many students find challenging is the relationship between acids, bases, and the pH scale. For the purposes of this article, the Arrhenius definition of acids and bases will apply. According to this definition, acids are any substance that gives off one or more protons (H+ ions) in solution whereas a base is anything that releases hydoxide (-OH) ions into solution.
Empirically, acids taste sour, and strong acids are generally corrosive. Familiar acids include vinegar (acetic acid), lemon juice (citric acid), HCl (hydrochloric acid), and H2SO4 (sulfuric acid) — present in car batteries. Even someone who has never taken chemistry will recognize that vinegar and lemon juice are much milder than battery acid. Still, a question naturally arises: what makes the latter two acids so much "stronger" than the first two?
The answer is that the more readily a compound gives off H+ ions, the stronger an acid it makes. Similarly, strongly alkaline or basic substances release many hydroxide ions into solution. Bases are also corrosive; however, unlike acids they taste bitter, not sour. Examples of strong bases include lye (used in soap production) as well as drain cleaners like Liquid Plumber or Drano. Weaker bases including sodium bicarbonate and calcium hydroxide are used in antacid tablets.
This phenomenon is a classic example of a double replacement reaction, known in acid-base chemistry as a neutralization reaction:
Acid + Base -> Salt + Water
The H+ ion from the acid and -OH ion from the base combine to form water while the remaining components form a salt. Table salt, better known as sodium chloride or NaCl, is formed by a neutralization reaction between hydrochloric acid and the base sodium hydroxide:
HCl + NaOH -> NaCl + H2O
Quantifying Acid and Base Strength: the pH scale
Once chemists could accurately measure the acidity or alkalinity of solutions, the next step was to establish a numeric scale based on logarithmic values, or powers of ten. There are several important concepts to keep in mind when learning the pH scale. If you remember nothing else from this article, learn these key points:
1) pH is defined as - log [H+].
2) Each whole number on the pH scale represents a ten fold difference in H+ concentration.
3) The scale runs from 0 to 14. This covers the spectrum of acids and bases commonly encountered in nature as well as most laboratories.
4) Strong acids have a low pH.
5) Strong bases have a high pH.
6) Pure water is neutral and has a corresponding pH of 7.
Water, for all intents and purposes, is neutral due to its extremely low dissociation value of 10 to the -7 power. This means out of every ten million water molecules only one spontaneously dissociates into H+ and -OH ions. Therefore, the pH of pure water is equal to 7, which can be thought of as the neutral midpoint of the pH scale.
Compared to water, strong acids like HCl have an extremely low pH, i.e. zero.
This is the part many people find confusing. A pH of 0 does NOT mean the absence of anything but rather that nearly every molecule in the solution has given off an H+ ion.
A solution with a pH of 0 contains 10 to the zero power of protons, which equals 1 mole of H+ ions per liter. This is a tremendously large number of protons (6.02 x 10 to the power of 23 to be exact).
In comparison, a mildly acidic solution like dilute vinegar (acetic acid) has a pH of around 4, which is one thousand times more acidic than pure water but ten times less acidic than lemon juice (pH ~ 3), and ten thousand times less acidic than battery acid. In other words, for every ten thousand molecules of acetic acid, just one of them has given off a proton, leaving behind a negatively charged acetate ion.
Mathematically, 1/10,000 is expressed as 10 to the -4 power. Since by definition pH = - log [H+]; substituting (10)-4 for the H+ concentration gives a pH of 4.
In contrast to acids, which become stronger as the pH decreases, bases are the opposite. They become stronger as the pH increases.
Another way to measure the strength of a base is to subtract its pH from 14. The difference is a value called pOH, which can be thought of as the basic or alkaline counterpart of pH. Strong bases have a low pOH; weak bases have a pOH of 6 to slightly less than 7.
For example, suppose a dissolved antacid tablet has a pH of 9. Only one free proton exists for every one billion molecules of antacid in solution. Subtracting pH 9 from pH 14 gives a measure of alkalinity of pOH = 5. In this case, one in one hundred thousand antacid molecules has released an -OH ion.
Compared to weak bases, strong bases like sodium hydroxide, found in oven cleaners like Easy Off, have a pH close to 14 (and a corresponding pOH near 0). They contain approximately 1 mole of -OH ions per liter of solution.
In addition to reacting violently with each other, concentrated solutions of acids and bases react with most metals, dissolving them in the process. As such, these highly reactive substances must be stored in chemically inert containers made of glass or certain plastics.