Definition of Acids and Bases

Need help preparing for the General Chemistry section of the MCAT? MedSchoolCoach expert, Ken Tao, will teach you the definitions of Acids and Bases. Watch this video to get all the MCAT study tips you need to do well on this section of the exam!
Acids and bases are an incredibly high yield topic for the MCAT, so we will cover their definitions, applications, and uses in great detail. Here we lay the groundwork for this topic by defining acids and bases, discussing their relationship with water, and contextualizing the pH and pOH scales.
Definitions of Acids and Bases
There are two primary definitions of acids and bases used by the MCAT: The Brønsted-Lowry definition and the Lewis definition. Under the Brønsted-Lowry definition, an acid is defined as a hydrogen ion (H+) donor. Remember, a hydrogen atom is a proton and an electron. If you take an electron away, the hydrogen ion is literally just a proton. Conversely, a Brønsted-Lowry base is a hydrogen ion acceptor.
Under the Lewis definition, an acid is an electron pair acceptor. Conversely, a Lewis base is an electron-pair donor. While the Lewis definition is more widely used in an organic chemistry context, for the purposes of our acid-base video notes, we will be considering acids to be hydrogen ion donors and bases as hydrogen ion acceptors (or, shorthand, as proton donors and acceptors, respectively).
The Auto-ionization of Water
The auto-ionization of water is a process by which pairs of water molecules spontaneously react with one another to form hydronium (H3O+) and hydroxide (OH-). The reaction can be described with the equation below, in which one molecule of water is acting as an acid and another as a base.
On the product side of the reaction, hydroxide is what we call the conjugate base. The definition of a conjugate base is the molecule that’s left over after the acid has donated its proton. The hydronium ion is what is called the conjugate acid. The conjugate acid is simply the base after it has accepted a proton. In this reaction, water is acting as both an acid and a base. When a molecule is acting as both an acid and a base, it is termed an amphoteric molecule.
Knowing the form of the reaction, the next question we may ask is, “How much ionization actually occurs?”. The equilibrium constant (the ratio of products to reactants at equilibrium) for the autoionization of water is Kw. The expression for Kw is below, notably missing is H2O from the denominator, as liquids and solids are excluded from equilibrium expressions. The value of Kw (which you must have memorized) is 1 x 10-14 at 25°C Like other equilibrium constants, the value will differ at different temperatures.
Note that the equilibrium expression for the autoionization of water suggests that hydronium and hydroxide are produced in equal amounts, the product of which equals 1 x 10-14. Therefore, in pure water at 25oC, the concentration of the hydronium ion and hydroxyl ions are both 10-7 M.
pH and pOH
The pH of a solution is a measure of the concentration of hydronium in a solution. pH can be calculated using the equation -log[H3O+]. Operating on a logarithmic scale, values of pH that differ by 1 imply that hydronium concentration differs 10-fold. For instance, a solution at pH 1 has 10 times as many hydronium ions as a solution at pH 2, and a solution at pH 2 has 10 times more hydronium than a solution at pH 3. The values of pH can generally range from 0 to 14. A pH value of seven is defined as a neutral solution - the concentration of hydronium and hydroxide are equal. A pH of less than 7 is acidic because the concentration of hydronium ion is greater than the concentration of hydroxide. A pH greater than 7 is basic, as the concentration of hydroxide is greater than the concentration of hydronium.
Conversely, the pOH of a solution is equal to -log[OH-]. The pOH also contains values ranging from 0 to 14, although a pOH below 7 is basic and vice versa.
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