Learn high-yield acids and bases concepts for the MCAT, including Bronsted–Lowry theory, Henderson–Hasselbalch, and practice questions.

(Note: This guide is part of our MCAT General Chemistry series.)

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.

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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

For the MCAT, it’s important to memorize several common strong acids and bases. Soluble hydroxides of group 1 and group 2 of the periodic table form strong bases when they are compounded with the hydroxide ion. The one exception to this is Be(OH)₂. Below is a list that summarizes several commonly found strong bases. Any bases not found on this list can be assumed to be weak.

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).

H₂O + HA ⇋ H₃O⁺ + A^-
H₂O + B^- ⇋ OH^- + 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:

H₂O + H₂O ⇋ H₃O⁺ + OH^-

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.)

K_w, or the water dissociation constant, is an important value to memorize. At 298 Kelvin (25°C), the K_w 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: K_a * K_b = K_w = 10^-14
Two additional constants to be familiar with are the acid dissociation constant (K_a) and the base dissociation constant (K_b). These are the equilibrium constants of the dissociation of an acid (HA) and the dissociation of a base (BH). The formulas for both K_a and K_b are shown below.

HA + H₂O ⇋ H₃O⁺ + A^- K_a = [H₃O⁺][A^-]/[HA]
B + H₂O ⇋ BH⁺ + OH^- K_b = [BH⁺][OH^-]/[B]
Since strong acids dissociate completely, strong acids typically have a larger K_a value. Weak acids have a K_a less than 1.0. Similarly, a larger K_b value corresponds with a strong base, and weak bases have a K_b less than 1.0. The negative logarithms (base 10) of the acid dissociation constant and base dissociation constants provide several important values. pK_a = -log(K_a)
pK_b = -log(K_b)
The pK_a 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 pK_a, the species will be mostly deprotonated. Strong acids tend to have low pK_a values, and strong bases tend to have low pK_b 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) (K_a = 6.46 x 10^-5).

First, write out the chemical reaction.

C₆H₅COOH + H₂O ⇋ C₆H₅COO^- + H₃O⁺
Using this, the K_a expression can be written: K_a = [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 K_a expression can then be rewritten as: K_a = (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:
K_a = (x)(x)/(3.0)
Lastly, the provided K_a value can be substituted and used to solve for x:

6.46 x 10^-5 = (x)(x) / (3.0)
0.0001938 = x²
x = 0.014
Thus, the concentration of hydronium in the solution is equal to 0.014 M.

e) pH and pOH

The pH scale is used to describe the acidity of a solution. A pH value below 7 is considered acidic, whereas a pH value above 7 is basic. A pH of 7 is neutral and is the pH of pure water. A strong acid will typically have a pH between 0 to 3, whereas a weak acid will have a pH between 3 and 7. A strong base will typically have a pH between 12 to 14, whereas a weak base will typically have a pH between 7 and 12.

The pH of a solution can be found using the following equation:

pH = -log[H⁺]

The pH is equal to the negative log of the concentration of hydrogen ions in a solution. A similar scale, the pOH scale, provides the basicity of a solution. It is essentially the opposite of the pH scale. A pOH below 7 is considered basic, above 7 is acidic, and at 7 it is neutral. The equation below can be used to find the pOH of a solution.

pOH = -log[OH^-]

The pOH is equal to the negative log of the concentration of hydroxide ions in a solution. For the same solution, its pOH and pH both add up to 14. For example, if a pH of a solution is measured to be 4. Then the pOH of that solution must be 10. This is summarized by the equation:

pH + pOH = 14

Determining the pH or pOH of a solution of a strong acid or strong base is fairly easy. Consider the following:

A 1L aqueous solution contains 0.001 M hydrochloric acid. Find its pH.

Since hydrochloric acid (HCl) is a strong acid, it dissociates completely in solution.

HCl + H₂O ⇋ Cl^- + H₃O⁺
Thus, [H₃O⁺] must also be equal to 0.001. This value can be used in the definition of pH:

pH = -log[H⁺]
pH = -log(0.001)
pH = -(-3) = 3
However, this becomes much more difficult for weak acids and bases. Since weak acids and bases do not completely dissociate in an aqueous solution, an alternative method must be used to determine the final pH.

If the pK_a of a weak acid or base is known, the concentration of H⁺ can be determined using the method described in the previous section. Then, this value can be used in the definition of pH.

Find the pH of a 3.0 M aqueous solution of the weak acid, benzoic acid (C₆H₅COOH) (K_a = 6.46 x 10^-5). Based on our previous calculations, we know that the concentration of hydronium in the solution is equal to 0.014 M. Then, pH = -log[H⁺]
pH = -log(0.014)
While this may be a difficult value to obtain without a calculator, the following is true:

-log(0.010) > -log(0.014) > -log(0.1)
-log(0.010) > pH > -log(0.1)
-(-2) > pH > -(-1)
2 > pH > 1
Thus, the pH of the solution is assumed to be between 1 and 2. (To be precise, the exact value of the pH is 1.85.)

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Part 3: Buffer systems

a) Buffer systems

Buffer systems contain a mixture of a weak acid or base and a salt containing its conjugate ion. As the salt of the conjugate ion is added, it undergoes hydrolysis and dissolves to form the weak acid or base.

Buffer systems are used for their ability to resist pH changes. Let’s examine a common buffer solution, the acetic acid and acetate buffer solution. Here, acetic acid is the weak acid, while acetate is the conjugate base.

CH₃COOH + H₂O ⇋ CH₃COO^- + H₃O⁺

If a strong base were to be added to the solution, it would react with water to form hydronium. To counteract this, acetic acid dissociates to form the acetate ion. Since acetate is a weak base, the pH changes very little compared to the strong base.

The bicarbonate buffer system is an extremely important buffer system within our bodies. This buffer system is used to regulate the pH of blood. For more information on the bicarbonate buffer system, be sure to refer to our guide on the respiratory and cardiovascular systems.

b) Henderson-Hasselbalch equation

The Henderson-Hasselbalch equation is used to determine the pH of a buffer system. The equation is written as:

pH = pK_a + log([A^-]/[HA])
where A^- represents the conjugate base and HA represents the protonated acid. Let’s apply this to an example problem.

A 1L aqueous solution contains 0.001 M acetic acid and 0.1 M acetate. Find its pH. (Note: The K_a of acetic acid is 1.8 x 10^-5).

First, write down the expression for the chemical reaction. Here, acetic acid dissociates to form acetate, its conjugate base.

CH₃COOH + H₂O ⇋ CH₃COO^- + H₃O⁺
The question stem states that the concentrations that must be used in the Henderson-Hasselbalch equation. Additionally, the K_a is provided. This can be converted to pK_a by taking the negative logarithm of the provided K_a.

pH = pK_a + log([A^-]/[HA]
pH = pK_a + log([CH₃COO^- ]/[CH₃COOH ]
pH = -log(1.8 x 10^-5) + log(0.1/0.001)
pH = 4.7 + log(100)
pH = 4.7 + 2
pH = 6.7
The buffer range is the effective range over which a buffer system is effective at resisting changes in pH. A buffer range can be determined by performing a titration of the weak acid or base.

Common Buffer Systems

    SAC329 code block
  

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Part 4: Titrations

a) Titrating a strong acid and strong base

Titrations allow us to find the concentration of a reactant in a solution. This is done by taking a small amount of a titrant, which is a solution of known concentration, and adding it to an analyte, which is a solution of unknown concentration. At a certain point, the number of moles of acid and the moles of the base will equalize. This occurs at the equivalence point.

At the equivalence point, the following equation can be used:

N_acidV_acid = N_baseV_base
where N is the normality of the acidic or basic solution,
V is the volume of acid or base used

If three of these variables are known, then the equation can be used to solve for the fourth.

The normality refers to the number of equivalents of acid or base that are used in a solution. For an acid with a single dissociating proton such as HCl, this is equal to the concentration (in moles) of the solution. For a multiprotic acid with multiple dissociating protons, this is roughly equal to the number of acidic protons times the concentration of the solution.

To find the endpoint or equivalence point of an acid-base titration, indicators are used. Indicators are often weak acids or bases themselves. These molecules change colors in response to acidity and basicity, thus providing a visual indication of a solution’s pH. To select an ideal indicator, an indicator with a pK_a value closest to the equivalence point pH should be used. Below is a list of a few indicators and their behaviors.

Table 1 Key indicators and their characteristics

Name of indicator	Range of color change	Color
Methyl red	pH 4-6	Orange below pH of 4.4; yellow at pH above 6.3
Bromothymol blue	pH 6-7.5	Green in neutral solution; yellow in acidic solution; blue in basic solution
Phenolphthalein	pH 8-10	Colorless at pH below 8; pink at pH above 9

Below is the titration curve of a strong base with a strong acid. Initially, the pH value is high as the solution is a strong base. However, as more strong acid is added, an equivalence point is achieved. The further addition of strong acid then results in a further decrease of pH.

**Figure** 1 Titrating a strong base with a strong acid

In a titration curve, the equivalence point can be identified as the midpoint of the region with the steepest slope. When a titration involves a strong acid and strong base, the equivalence point occurs at a pH of 7.

The exact opposite trend should be expected in titrating a strong acid with a strong base: the pH of the solution will start at a low value, achieve an equivalence point at pH = 7, then continue to rise.

The equivalence point of a titration involving a strong acid and strong base always occurs at a pH of 7. When this point is reached, neutralization is said to have occurred: the solution is no longer acidic or basic.

b) Titrating weak acids and bases

When titrating a strong acid with a weak base, the equivalence point will be at a pH that is below 7. Thus, the final solution will still be fairly acidic. However, when a strong base is titrated with a weak acid, the equivalence point will be at a pH that is above 7. Thus, the final solution will still be fairly basic.

Below is the titration curve of a strong base titrated with a weak acid. The pH starts off at just below 12, as the initial solution is strongly basic. However, as more weak acid is added, an equivalence point above 7 is reached.

**Figure 2** A strong base titrated with a weak acid.

Below is the titration curve of a weak acid being titrated with a strong base. The pH starts off at around 4 since the initial solution is weakly acidic. However, as more of the strong base is added, an equivalence point just above 7 is achieved. The pH then continues to rise as more base is added.

**Figure 3** A weak acid titrated with a strong base.

c) pKa of multiprotic acids

In a multiprotic acid, each proton has a different pK_a. This results in multiple equivalence points and midpoints. The midpoint is the point at which half of the sample is 50% deprotonated and 50% protonated. A midpoint is also known as a half-equivalence point. It can be estimated as the point at the center of the portion of the curve with the flattest slope. The figure below identifies the half-equivalence and equivalence points of a multiprotic titration curve.

**Figure 4** Titration of a multiprotic acid

The regions of low slope in which half-equivalence points are found are also known as buffer regions. Recall that this buffer region is the range of pH values over which the acid or base and its conjugate ion are effective as a buffer system.

If several pK_a values are known, they can be used to find the isoelectric point, or pI. This is the pH at which a compound, such as an amino acid, is completely neutrally charged.

The average of a compound’s pK_a values provides its pI. For instance, if pK_a1 of a compound is equal to 3 and pK_a2 is equal to 4, the pI is equal to 3.5.

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Part 5: High-yield terms

Amphoteric species: species that act as either acids or bases

Analyte: reactant that is of unknown concentration in a titration

Autoionization: a process by which water reacts with itself

Bronsted-Lowry acid: species that donates a proton

Bronsted-Lowry base: species that accepts a proton

Buffer solutions: composed of a weak acid/base and a salt containing its conjugate; used for its ability to resist pH changes

Conjugate acid: a base that has acquired a hydrogen ion

Conjugate base: an acid that has lost a proton

Equivalence point: occurs when the number of equivalent acids and bases are equalized in a titration

Indicator: a molecule that changes color in response to acidity/basicity; determines the endpoint of a titration

Lewis acid: species that accepts an electron pair

Lewis base: species that donates an electron pair

Midpoint: occurs when the concentration of deprotonated species is equal to the concentration of protonated species

pH scale: describes the concentration of hydrogen ions in a solution

pOH scale: describes the concentration of hydroxide ions in a solution

Titrant: reactant which is of known concentration in a titration

Titration: a process that allows us to find the concentration of a reactant in a solution

Zwitterions: molecules that have both a positive and negative charge

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Part 6: Passage-based questions and answers

The strength of an acid is determined by its tendency to dissociate into a proton and a conjugate base. A strong acid is one that dissociates completely in a solution.

HA → H⁺ + A^-
Reaction 1: The dissociation of a generic strong acid
In contrast to a strong acid, a weak acid only partially dissociates. The undissociated reactants and the dissociated products remain in an equilibrium state.

HA ⇋ H⁺ + A^-
Reaction 2: The dissociation of a generic weak acid
Ethanoic acid, commonly known as acetic acid, is an important industrial chemical reagent. It is found in many household substances, ranging from descaling agents to condiments. Several students investigate the dissociation of ethanoic acid in an aqueous solution.

CH₃COOH + H₂O ⇋ CH₃COO^- + H₃O⁺
Reaction 3: The dissociation of ethanoic acid
Additionally, the students observe the reaction between ethanoic acid and sodium hydroxide, a strong base.

CH₃COOH + NaOH ⇋ NaCH₃COO^- + H₂O
Reaction 4: Ethanoic acid and sodium hydroxide

This passage was created for educational purposes only

Question 1: Which of the following is not a strong acid?

A) Hydrochloric acid

B) Hydroiodic acid

C) Hydrofluoric acid

D) Sulfuric acid

Question 2: Which of the following accurately describes ethanoic acid?

A) Weak acid

B) Strong acid

C) Titration indicator

D) Conjugate base

Question 3: What is the concentration of hydronium in a 2.0 M aqueous solution of ethanoic acid? (Note: K_a = 1.8 x 10^-5)

A) 2 x 10^-3

B) 6 x 10^-3

C) 7 x 10^-3

D) 9 x 10^-3

Question 4: The concentration of the base used in titrating ethanoic acid is 1 x 10^-5 M. What is the pH of the solution?

A) 3

B) 5

C) 7

D) 9

Answer key for passage-based questions

Answer choice C is correct. Most halogens form strong acids. The exception to this is fluorine, since hydrofluoric acid is not considered a strong acid (choice C is correct). The strong acids you should know for the MCAT include hydroiodic acid, hydrobromic acid, hydrochloric acid, perchloric acid, sulfuric acid, and nitric acid (choices A, B, and D are incorrect).
Answer choice A is correct. Ethanoic acid, as the name suggests, is an acidic compound. The removal of a proton from ethanoic acid results in the formation of a conjugate base (choice C is incorrect). In this context, ethanoic acid is not used as a titration indicator (choice D is incorrect). Based on the information provided in reaction 3, ethanoic acid does not completely dissociate in solution and instead exists in an equilibrium state consisting of its undissociated reactants and dissociated products. Thus, it is a weak acid (choice B is incorrect).
Answer choice B is correct. First, identify the Ka expression using reaction 3. The provided value of Ka can be used.

K_a = [H₃O⁺][CH₃COO^-]/[CH₃COOH]
1.8 x 10^-5 = [H₃O⁺][CH₃COO^-]/[CH₃COOH]

Next, use x to express the amount of ethanoic acid and hydronium formed. These two quantities must be equal, since their stoichiometric ratios are 1:1. Since the amount of products formed will be equal to the initial concentration of ethanoic acid minus x, the K_a expression can be written as:

1.8 x 10^-5 = (x)(x)/(2.0 - x)

Since the value of x value is very small, it is negligible. Thus, this expression can be approximated as:

1.8 x 10^-5 = (x)(x)/(2.0)

Finally, solve for x.

1.8 x 10^-5 = (x)(x)/(2.0)
3.6 x 10^-5 = x²
6 x 10^-3 = x

Thus, the concentration of hydronium present is 6 x 10^-3 M.

4. Answer choice D is correct. First, use the hydroxide concentration to calculate the pOH value:

pOH = -log[OH^-]
pOH = -log(1 x 10^-5)
pOH = 5

To find the pH, simply subtract the pOH from 14.

pH + pOH = 14
pH = 14 - pOH
pH = 14 - 5
pH = 9

Thus, the pH of the solution is 9.

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Part 7: Standalone questions and answers

Question 1: Which of the following accurately characterizes a Lewis acid?

A) A Lewis acid is a species that accepts an electron pair

B) A Lewis acid is a species that donates an electron pair

C) A Lewis acid is a species that donates a proton

D) A Lewis acid is one that accepts a proton

Question 2: Which of the following is not a strong base?

A) Be(OH)₂

B) LiOH

C) CsOH

D) Mg(OH)₂

Question 3: What are amphoteric species?

A) Species that contain both hydrophilic and hydrophobic groups

B) Species that can act as either acids or bases

C) A species that contains both positive and negative charge

D) A species that is lacking a charge

Question 4: A phosphate buffer system contains 8 M H₂PO₄^- and 0.8 M HPO₄^2-. What is the pH? (Note: pK_a = 6.8)

A) -1

B) 5.8

C) 8

D) 8.8

Question 5: What is the isoelectric point of glycine?

A) 2.0

B) 4.5

C) 5.5

D) 9.0

Answer key for standalone questions

Answer choice A is correct. A Lewis acid is any chemical species that accepts an electron pair (choice A is correct). A Lewis base is any chemical species that donates an electron pair (choice B is incorrect). Bronsted-Lowry acids donate a proton, while Bronsted-Lowry bases accept a proton (choices C and D are incorrect).
Answer choice A is correct. Soluble hydroxides of group IA and IIA metals form strong bases (choices B, C, and D are incorrect). The one exception to this, however, is Be(OH)2.
Answer choice B is correct. Amphoteric species are those that are able to act as either acids or bases. One such example is water (choice B is correct). Amphipathic species are those that contain hydrophilic and hydrophobic groups (choices A is incorrect). Zwitterions are those that contain both a positive and negative charge (choice C is incorrect). A species that is lacking a charge is neutral (choice D is incorrect).
Answer choice B is correct. To solve this problem, the provided values can be used in the Henderson-Hasselbalch equation.

Note that H₂PO₄^- is the acid, while HPO₄^2- is the conjugate base.

pH = pK_a + log([A^-]/[HA])
pH = 6.8 + log(0.8/8)
pH = 6.8 + log(0.1)
pH = 6.8 + (-1)
pH = 5.8

Thus, the pH of the buffer system is 5.8.
5. Answer choice C is correct. Glycine, like all amino acids, has a carboxyl group (pK_a = 2) and amino group (pK_a = 9). To find the isoelectric point (pI), simply take the average of these two values.

pI = (pK_{a, 1} + pK_{a, 2}) / 2
pI = (2 + 9)/2
pI = 5.5

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Acids and Bases for the MCAT: Everything You Need to Know

Learn high-yield acids and bases concepts for the MCAT, including Bronsted–Lowry theory, Henderson–Hasselbalch, and practice questions.

Table of Contents

Part 1: Introduction to acids and bases