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a) Buffer systems
3 1/2 Approaches To MCAT Acids And Bases
Acid-base chemistry has long been, and still is, a high-yield topic on the MCAT. In chemistry, there are three approaches to classifying acids and bases. Each approach has its own weight on the exam and tends to appear more often in one section. In this blog, we will introduce all three, but focus on the two highest-yield areas.
Approach 1 – Arrhenius Theory
The most specific definition of acids and bases is the Arrhenius approach. Arrhenius theory posits that in order for a substance to be an acid or a base, it must release either H + or OH – ions (i.e. the compound must contain that ion). This definition covers many of the commonly known acids and bases (H 2 SO 4 , HNO 3 , Ca(OH) 2 , NaOH, KOH, etc.). However, this approach limits acidic and basic species to the aqueous phase. An Arrhenius acid is a compound that releases H + ions when added to water, and an Arrhenius base is a compound that releases OH – ions when added to water. Arrhenius theory is much more limited than the other two theories so it is the least tested on the MCAT. For example, Arrhenius theory fails to explain the basic behavior of amines (e.g. ammonia, NH 3 ) which, in the presence of water, increase hydroxide ions concentration in solution, but do not contain any OH – ions. This leads to the next, broader definition of acids and bases.
Approach 2 – Bronsted-Lowry Theory
The Bronsted-Lowry theory posits that an acid is any species which is a proton (H + ion) donor and a base is any species which is a proton acceptor. This includes species like ammonia and water. Water itself can act as both a BL-acid and a BL-base, making it an amphoteric species. A basic salt, such as Ca 2 +( F – ) 2 , will generate OH – ions in water by taking protons from water itself (to make HF). This approach forms the basis for pH, pOH, K a , K b , and K w . These values are tested quite often in the areas of general chemistry and biochemistry. An MCAT favorite in this area is to ask students to recognize the charges that can appear or disappear on amino acids and their side chains. The relative strengths of acid/base species is also based on the Bronsted-Lowry theory. This content area is math-heavy, so be sure to practice your calculations, work on your test day scratch work, and utilize your Next Step materials to learn valuable shortcuts to avoid math whenever possible.
Approach 3 – Lewis Theory
The most general school of thought on amino acids is the Lewis theory. This theory does not define acids or bases based on ions or proton exchange, but on the transfer of electrons between species. A Lewis acid is an electron pair acceptor while a Lewis base is an electron pair donor . Lewis bases (e.g. NH 3 ) are generally the same molecules seen in the Bronsted-Lowry definition but Lewis acids are different. Since Lewis theory posits that an acid species must be able to accepting an electron pair, many are species like H + ions (protons) that are cations. Many Lewis acids will have a positive charge and an empty orbital that can hold the electron pair. Other Lewis acids need not be ionic at all. Classic MCAT Lewis acids include BF 3 and AlCl 3 , the latter of which is a common catalyst for electrophilic aromatic substitution reactions (aromatics are fair game for the exam). This definition of acid base chemistry tends to show up more in the organic chemistry concepts, but may be combined with biochemistry. Keeping track of electrons is a big part of understanding the biochemistry tested on the MCAT.
Bonus – Schiff Base
A creative approach by the new MCAT is the concept of Schiff bases. With carbonyl derivatives, the AAMC has long tested the idea that the electrophilic carbonyl carbon is a ripe target for nucleophilic attack. If the attacking nucleophile is an amine, and the electrophile is an aldehyde or ketone, the product of this reaction is a compound in which the C=O double bond (carbonyl) is replaced by a C=N double bond (imine). This imine is also known as a Schiff base . Schiff base formation has important biological applications to metabolism, aldol reactions, antibiotic mechanisms, and the reactions of B vitamins. This concept is most likely to appear in Biological/Biochemical Foundations passages, but can also be tested under the guise of organic chemistry basics (e.g. nucleophilic attack).
No matter the school of thought, the test makers will expect you to recognize the actions of acids and bases in many contexts, so be sure to practice plenty of MCAT-style questions once you think you have a handle on the content. Looking to get more practice? Our online Qbank contains over 1,000 discrete MCAT science questions plus CARS practice passages! Add the Qbank to your NextStepMCAT.com account here.
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Acids and Bases for the MCAT: Everything You Need to Know
Learn key MCAT concepts about acids and bases, plus practice questions and answers
(Note: This guide is part of our MCAT General Chemistry series.)
Table of Contents
Part 1: Introduction to acids and bases
Part 2: Definitions
a) Lewis and Bronsted-Lowry acids and bases
b) Important acids and bases
c) Amphoteric species
d) Equilibrium constants
e) pH and pOH
Part 3: Buffer systems
a) Buffer systems
b) Henderson-Hasselbalch equation
Part 4: Titrations
a) Titrating a strong acid and strong base
b) Titrating weak acids and bases
c) pKa of multiprotic acids
Part 5: High-yield terms
Part 6: Passage-based questions and answers
Part 7: Standalone questions and answers
Part 1: Introduction to acids and bases
Maintaining equilibrium in acids and bases is a key component of maintaining homeostasis. Understanding the behavior of acidic and basic environments is key to understanding the basis of many organic reactions and biological systems.
This guide will cover properties that define acids and bases, buffer systems, and the fundamentals of titrations. By the end of this guide, you’ll be prepared for whatever the MCAT throws at you on acids and bases come test day. At the end of this article, there are also several MCAT-style practice questions for you to test your knowledge with.
Let’s get started!
Part 2: Definitions
a) Lewis and Bronsted-Lowry acids and bases
Many students think of acids and bases in terms of hydrogen and hydroxide ions. For example, a solution that has a high concentration of hydrogen ions (H + ) is considered acidic, while a solution that has a high concentration of hydroxide ions (OH – ) is considered basic. While this is certainly a valid way of looking at acids and bases, it’s not the only one!
The Lewis and Bronsted-Lowry definitions describe different properties of acids and bases. Lewis acids and bases are based on electron movement. A Lewis acid accepts an electron pair, while a Lewis base donates a pair of electrons.
Out of all the definitions of acids and bases, the Lewis definitions are the most inclusive. The Bronsted-Lowry definitions are slightly less inclusive and focus on proton movement. A Bronsted-Lowry acid is one that donates a proton (H + ), while a Bronsted-Lowry base is one that accepts a proton. Every Bronsted-Lowry base is also a Lewis base. However, not every Lewis base is a Bronsted-Lowry base.
All acids and bases have conjugates, which are formed after losing or gaining a hydrogen ion. A conjugate acid is a Bronsted-Lowry base that has acquired a proton. A conjugate base is a Bronsted-Lowry acid that has lost a proton. Strong acids have weak conjugate bases, and strong bases have weak conjugate acids. This is because a strong acid will not form a species that readily accepts a proton. If it did, then it would not be a strong acid.
b) Important acids and bases
- Potassium hydride: KH
- Lithium hydride: LiH
- Lithium hydroxide: LiOH
- Sodium hydroxide: NaOH
- Potassium hydroxide: KOH
- Rubidium hydroxide: RbOH
- Caesium hydroxide: CsOH
- Magnesium hydroxide: Mg(OH)2
- Calcium hydroxide: Ca(OH)2
- Strontium hydroxide: Sr(OH)2
- Barium hydroxide: Ba(OH)2
Elements from Group 17, the halogens, typically form strong acids. These compounds form very stable conjugate bases upon dissociation due to the large atomic radius of the resulting ion. Hydrobromic acid (HBr) dissociates into a hydrogen ion and bromide ion. The bromide ion’s larger size gives the negative charge more room to disperse, thus stabilizing the ion. The exception to this rule is hydrofluoric acid (HF). The weakness of HF is a result of the instability of the dissociated products. The negative charge of the fluoride anion is very unstable due to the small atomic radius of the ion. Thus, the dissociation of HF is not very favorable. The list below summarizes several commonly found strong acids. Any acids not on this list can be assumed to be weak.
- Hydroiodic acid: HI
- Hydrobromic acid: HBr
- Hydrochloric acid: HCl
- Perchloric acid: HClO4
- Sulfuric acid: H2SO4
- Nitric acid: HNO3
c) Amphoteric species
Amphoteric species are compounds that can act as either Bronsted-Lowry acids or bases. In an acidic environment, such a species acts as a base, while in a basic environment, it acts as an acid. Amino acids are an example of amphoteric species. Recall that amino acids are classified as zwitterions. These are molecules that simultaneously have a positive and negative charge. Amino acids can act as acids through the amino group (or N-terminus), which donates protons. Additionally, they can act as bases through their carboxylate group (or C-terminus), which accepts protons.
Water is another very common amphoteric species. It undergoes different reactions when in the presence of a base or acid. Below are examples of water in the presence of a protonated acid (HA) and a protonated base (HB).
Note that when water is in an acidic environment, it accepts a proton to form the hydronium ion. However, when it is in a basic environment, it donates a proton to form a hydroxide ion. Water is also able to engage in a special type of reaction called autoionization. During this process, water reacts with itself. This is outlined by the following reaction:
d) Equilibrium constants
The autoionization of water also has an equilibrium constant. (For more information on equilibrium constants, be sure to refer to our guide on chemical equilibrium and kinetics.)
Kw, or the water dissociation constant, is an important value to memorize. At 298 Kelvin (25°C), the Kw is equal to the product of the concentrations of hydroxide and hydrogen ions, 10 -14 . Like all equilibrium constants, this is dependent only on temperature. Changes in factors such as concentration or volume will not have an effect on the water dissociation constant. The water dissociation constant can be used to calculate the acid dissociation constant or base dissociation constant of a solution: Ka * Kb = Kw = 10 -14
Two additional constants to be familiar with are the acid dissociation constant (Ka) and the base dissociation constant (Kb). These are the equilibrium constants of the dissociation of an acid (HA) and the dissociation of a base (BH). The formulas for both Ka and Kb are shown below.
HA + H2O ⇋ H3O + + A – Ka = [H3O + ][A – ]/[HA]
B + H2O ⇋ BH + + OH – Kb = [BH + ][OH – ]/[B]
Since strong acids dissociate completely, strong acids typically have a larger Ka value. Weak acids have a Ka less than 1.0. Similarly, a larger Kb value corresponds with a strong base, and weak bases have a Kb less than 1.0. The negative logarithms (base 10) of the acid dissociation constant and base dissociation constants provide several important values. pKa = -log(Ka)
pKb = -log(Kb)
The pKa is the pH at which half of the species in a solution are protonated, and the other half are deprotonated. If the pH is higher than the pKa, the species will be mostly deprotonated. Strong acids tend to have low pKa values, and strong bases tend to have low pKb values. Let’s take a look at an example.
Find the concentration of hydronium in a 3.0 M aqueous solution of the weak acid, benzoic acid (C6H5COOH) (Ka = 6.46 x 10 -5 ).
First, write out the chemical reaction.
C6H5COOH + H2O ⇋ C6H5COO – + H3O +
Using this, the Ka expression can be written: Ka = [H3O + ][C6H5COO – ]/[C6H5COOH] = 6.46 x 10 -5 How do we solve for [H3O + ]? Note that [C6H5COO – ] must be equal to [H3O + ] since their stoichiometric ratios are 1:1. We can use the variable x to express the amount of C6H5COO – and H3O + are formed. Further, the amount of products formed will be equal to the initial concentration of benzoic acid minus x. The Ka expression can then be rewritten as: Ka = (x)(x)/(3.0 – x)
Solving this equation is very difficult. To make the equation solvable, a simplifying assumption must be made. Since the value of x value is very small, it is negligible. Thus, the expression can be rewritten as: Ka = (x)(x)/(3.0)
Lastly, the provided Ka value can be substituted and used to solve for x:
6.46 x 10 -5 = (x)(x) / (3.0)
0.0001938 = x 2
x = 0.014
Thus, the concentration of hydronium in the solution is equal to 0.014 M.
