Unit 8: Acids and Bases

This unit examines the properties, definitions, and reactions of acids and bases, along with how they behave in aqueous equilibrium. Students will learn how to calculate pH and pOH, determine acid and base strengths, and apply equilibrium concepts to weak acid/base systems and buffers. The unit also covers titration curves, polyprotic acids, and indicators, connecting these concepts to practical laboratory analysis and real-world chemical systems.

Introduction to Acids and Bases

The Arrhenius definition states that acids produce \( \text{H}^+ \) (protons) in aqueous solution, while bases produce \( \text{OH}^- \) ions. The Brønsted–Lowry definition is broader: acids donate protons and bases accept protons, which includes reactions that do not produce \( \text{OH}^- \) directly. The Lewis definition is even more general, defining acids as electron-pair acceptors and bases as electron-pair donors.
Strong acids and bases dissociate completely in water, meaning their equilibrium lies far to the product side, and \(K_a\) or \(K_b\) is very large. Weak acids and bases only partially dissociate, resulting in an equilibrium between undissociated and dissociated forms; these have much smaller \(K_a\) or \(K_b\) values. This distinction is essential for deciding whether to use stoichiometry alone (strong) or equilibrium calculations (weak).
The pH scale measures the acidity or basicity of a solution, defined as \( \text{pH} = -\log[\text{H}^+] \), with acidic solutions having pH < 7 and basic solutions having pH > 7 at 25°C. pOH is related by \( \text{pOH} = -\log[\text{OH}^-] \), and pH and pOH are connected through \( \text{pH} + \text{pOH} = 14 \) at 25°C. Knowing one value allows calculation of the other, making these relationships central to acid-base problem solving.
Water undergoes autoionization, producing small amounts of \(\text{H}^+\) and \(\text{OH}^-\) in pure form. The equilibrium constant for this process is \(K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}\) at 25°C. This constant is temperature-dependent, and any factor that changes \(K_w\) will also change the neutral pH value from 7.
Acid-base strength and reactivity are also influenced by molecular structure, electronegativity, and bond polarity. For binary acids, strength increases down a group and across a period due to bond strength and polarity effects. For oxyacids, strength increases with the number of oxygen atoms attached to the central atom, as more oxygen atoms draw electron density away from the O–H bond, making proton release easier.

Strong vs. Weak Acids and Bases

Strong acids dissociate completely in water, producing a high concentration of \(\text{H}^+\) ions and leaving virtually no undissociated acid molecules at equilibrium. Examples include HCl, HBr, HI, HNO₃, H₂SO₄ (first proton), and HClO₄. Because the dissociation goes essentially to completion, strong acids have very large \(K_a\) values, making equilibrium calculations unnecessary—pH is determined directly from initial concentration.
Weak acids only partially dissociate in water, producing an equilibrium mixture of undissociated acid and its conjugate base. Common examples are acetic acid (\(\text{CH}_3\text{COOH}\)), HF, and many organic acids. Weak acids have small \(K_a\) values, requiring equilibrium calculations to determine \([\text{H}^+]\) and pH, often using the ICE table method.
Strong bases dissociate completely in water to produce hydroxide ions, such as NaOH, KOH, and Ba(OH)₂ (which produces two \(\text{OH}^-\) per formula unit). These are typically soluble metal hydroxides from Group 1A and some Group 2A metals. Their pOH (and thus pH) can be calculated directly from the initial concentration of base.
Weak bases only partially react with water to produce hydroxide ions. Examples include ammonia (\(\text{NH}_3\)) and amines like methylamine (\(\text{CH}_3\text{NH}_2\)). These have small \(K_b\) values, meaning an equilibrium calculation is needed to determine hydroxide concentration and pH.
The strength of an acid or base refers to the extent of dissociation, not the concentration. A dilute solution of a strong acid can have a lower \([\text{H}^+]\) than a concentrated solution of a weak acid, so strength and concentration must not be confused. On the AP exam, this distinction is often tested with scenarios comparing solutions of different molarities and dissociation tendencies.

Acid and Base Dissociation Constants (\(K_a\) and \(K_b\))

The acid dissociation constant \(K_a\) measures the strength of a weak acid, quantifying the equilibrium between the undissociated acid and its ions in water. For a weak acid \(HA\), \(K_a = \frac{[H^+][A^-]}{[HA]}\), where all concentrations are equilibrium values. Larger \(K_a\) means greater dissociation and a stronger acid.
The base dissociation constant \(K_b\) measures the strength of a weak base, quantifying its equilibrium with water to produce \(\text{OH}^-\). For a weak base \(B\), \(K_b = \frac{[BH^+][OH^-]}{[B]}\). Larger \(K_b\) indicates greater production of hydroxide ions and a stronger base.
\(K_a\) and \(K_b\) are related through the water ionization constant: \(K_a \times K_b = K_w = 1.0 \times 10^{-14}\) at 25°C. This means knowing one constant allows calculation of the other for conjugate acid-base pairs. This relationship is frequently tested in multi-step AP exam problems involving conjugates.
For polyprotic acids (acids that can donate more than one proton), each proton dissociation has its own \(K_a\) value, with \(K_{a1} > K_{a2} > K_{a3}\) because it becomes progressively harder to remove additional protons. Often, only the first dissociation is considered in pH calculations because it dominates \([\text{H}^+]\) in solution.
On the AP exam, identifying whether to use \(K_a\) or \(K_b\) depends on the species given. If a problem gives a base but provides \(K_a\) for its conjugate acid, the relationship \(K_a \times K_b = K_w\) must be used to find the missing value before performing pH or pOH calculations.

The Relationship Between \(K_a\), \(K_b\), and \(K_w\)

For any conjugate acid–base pair, the product of their equilibrium constants equals the ionization constant of water: \(K_a \times K_b = K_w\). At 25°C, \(K_w = 1.0 \times 10^{-14}\), so knowing either \(K_a\) or \(K_b\) allows calculation of the other. This relationship applies only to conjugate pairs, such as \(\text{NH}_4^+\) (acid) and \(\text{NH}_3\) (base).
Stronger acids have larger \(K_a\) values, which means their conjugate bases are weaker (smaller \(K_b\)). Conversely, stronger bases have larger \(K_b\) values and weaker conjugate acids with smaller \(K_a\). This inverse relationship is a direct consequence of the fixed value of \(K_w\) at a given temperature.
This relationship is especially important when a problem gives you the equilibrium constant for one form (acid or base) but asks you to calculate properties of its conjugate. For example, if \(K_a\) of acetic acid is known, \(K_b\) for acetate can be found using \(K_b = \frac{K_w}{K_a}\). This step is often needed before calculating pH or pOH.
Temperature changes affect \(K_w\), which in turn changes the numerical values of \(K_a\) and \(K_b\) for conjugate pairs. While \(K_a \times K_b\) will still equal the new \(K_w\), the neutral pH value will shift from 7 if \(K_w\) changes. This is why pure water at higher temperatures has a pH less than 7 but is still neutral in terms of \([\text{H}^+] = [\text{OH}^-]\).
Remember that \(K_a\) and \(K_b\) are unitless when concentrations are expressed in molarity. On the AP exam, calculations using this relationship typically test your understanding of acid–base strength trends and your ability to move between acid and base properties mathematically.

pH and pKa

pKa is the negative logarithm of the acid dissociation constant: \(\text{p}K_a = -\log K_a\). A smaller pKa indicates a stronger acid because a larger \(K_a\) means more dissociation into \(\text{H}^+\) ions. The pKa scale is useful for quickly comparing relative acid strengths without working with very small decimal values for \(K_a\).
pKa values directly relate to pH in buffer solutions through the Henderson–Hasselbalch equation: \(\text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}\). When \([\text{A}^-] = [\text{HA}]\), the ratio is 1 and \(\text{pH} = \text{p}K_a\). This condition is important because it marks the point of maximum buffer capacity for a conjugate acid–base pair.
Knowing pKa helps predict the predominant form of a molecule at a given pH. If the pH is lower than the pKa, the acid form (\(\text{HA}\)) dominates; if the pH is higher, the base form (\(\text{A}^-\)) dominates. This principle is heavily used in biochemistry and pharmaceuticals to predict solubility and membrane permeability.
pKa values also provide insight into molecular structure and reactivity. Acids with electronegative atoms or resonance stabilization of the conjugate base have lower pKa values, reflecting their increased stability after losing a proton. Comparing pKa values between related molecules helps explain trends in acid strength.
On the AP exam, pKa problems often require converting between \(K_a\) and pKa, identifying the stronger of two acids, or applying pKa to buffer calculations. Being comfortable moving between these forms and interpreting them in context is essential for solving multi-step equilibrium problems.

pH and pOH Calculations

pH is defined as \(\text{pH} = -\log[\text{H}^+]\) and pOH as \(\text{pOH} = -\log[\text{OH}^-]\). At 25°C, the relationship \(\text{pH} + \text{pOH} = 14.00\) comes from \(K_w = 1.0 \times 10^{-14}\). Knowing either \([\text{H}^+]\) or \([\text{OH}^-]\) allows calculation of both pH and pOH.
For strong acids and bases, \([\text{H}^+]\) or \([\text{OH}^-]\) is found directly from the initial concentration because they dissociate completely. If the acid or base has multiple ionizable protons or hydroxides (e.g., H₂SO₄, Ba(OH)₂), multiply the molarity by the number of ions produced per formula unit. This makes strong acid/base pH calculations straightforward.
For weak acids and bases, pH and pOH require equilibrium calculations using \(K_a\) or \(K_b\) and an ICE table. The approximation \([HA]_{\text{initial}} - x \approx [HA]_{\text{initial}}\) can be used when the percent ionization is small (usually less than 5%). This assumption should be verified after solving for \(x\) to ensure accuracy.
For solutions containing both acids and bases, determine which species is stronger and dominates the pH. In mixtures with strong and weak acids or bases, the strong species sets the pH because it completely dissociates, while the weak species’ contribution is negligible. For salts of weak acids or bases, hydrolysis must be considered in pH calculation.
pH and pOH calculations are temperature-dependent because \(K_w\) changes with temperature. At higher temperatures, \(K_w\) increases, lowering the neutral pH below 7, while at lower temperatures, \(K_w\) decreases, raising the neutral pH above 7. On the AP exam, recognizing that pH 7 is not always “neutral” is a common point tested.

Percent Ionization

Percent ionization measures the fraction of a weak acid or base that dissociates in solution and is calculated as: \[ \% \text{Ionization} = \frac{[\text{H}^+]_{\text{equilibrium}}}{[HA]_{\text{initial}}} \times 100\% \] for acids, or \(\frac{[\text{OH}^-]_{\text{equilibrium}}}{[B]_{\text{initial}}} \times 100\%\) for bases. It provides a way to compare strengths of weak acids/bases of the same initial concentration.
For weak acids, a higher percent ionization means the acid is stronger because it produces more \(\text{H}^+\) ions at equilibrium. For weak bases, a higher percent ionization means more \(\text{OH}^-\) is produced, indicating greater base strength. Comparing percent ionization values at the same concentration avoids misleading conclusions caused by concentration differences.
Percent ionization increases as the initial concentration of a weak acid decreases. This is because the equilibrium shift is more significant when fewer acid molecules are present initially, even though the absolute \(\text{H}^+\) concentration is smaller. This trend is important for interpreting experimental data on weak acids and bases.
When calculating pH for weak acids or bases, percent ionization can be used to check whether the “\(x\) is small” approximation is valid. If the percent ionization is less than 5%, the change in concentration is small enough to justify ignoring \(x\) in subtraction. If it’s greater than 5%, the exact quadratic solution is preferred for accuracy.
On the AP exam, percent ionization questions often require you to explain results in terms of equilibrium shifts and the magnitude of \(K_a\) or \(K_b\). Understanding the connection between percent ionization, dissociation constants, and pH ensures accurate reasoning in free-response explanations.

Polyprotic Acids

Polyprotic acids can donate more than one proton per molecule, releasing them in sequential ionization steps. Each step has its own acid dissociation constant (\(K_{a1}\), \(K_{a2}\), \(K_{a3}\)), with \(K_{a1} > K_{a2} > K_{a3}\) because removing additional protons from an increasingly negative species is less favorable. Common examples include H₂SO₄, H₂CO₃, and H₃PO₄.
In calculations, the first dissociation usually determines the pH because \(K_{a1}\) is significantly larger than the subsequent constants. The contribution from later dissociations to \([\text{H}^+]\) is generally negligible unless very high accuracy is required or the acid is extremely dilute. For example, in H₂SO₄, the first proton dissociates completely while the second has a measurable but smaller \(K_a\).
Polyprotic acids can form multiple conjugate bases, each capable of acting as a weak acid or base in subsequent equilibria. For instance, H₃PO₄ can lose one proton to form H₂PO₄⁻, which can then lose another to form HPO₄²⁻, and finally PO₄³⁻. Each conjugate base’s behavior depends on the corresponding \(K_a\) or \(K_b\).
Polyprotic bases, such as carbonate (\(\text{CO}_3^{2-}\)), work in the reverse way, accepting protons in steps with decreasing \(K_b\) values. The multi-step equilibria for polyprotic systems can overlap with buffer behavior if a conjugate acid-base pair is present in significant amounts. This makes polyprotic species important in natural buffer systems, such as the bicarbonate buffer in blood.
On the AP exam, problems involving polyprotic acids often require identifying which dissociation step dominates pH, predicting the species present at a given pH, or calculating the pH of a solution containing a conjugate base. Clear understanding of sequential equilibria and relative \(K_a\) values is essential for these problems.

Molecular Structure and Acid Strength

Acid strength is influenced by the polarity and strength of the bond to the ionizable hydrogen. In binary acids (HX), strength increases down a group because bond length increases and bond strength decreases, making it easier to release \( \text{H}^+ \). Across a period, acid strength increases with electronegativity of X because a more polar H–X bond facilitates proton dissociation.
For oxyacids (\(\text{H–O–Y}\)), acid strength increases with the electronegativity of the central atom Y and with the number of oxygen atoms attached. Highly electronegative central atoms draw electron density away from the O–H bond, weakening it and making proton loss easier. Additional oxygens increase this effect through inductive electron withdrawal.
Resonance stabilization of the conjugate base also increases acid strength. If the negative charge on the conjugate base can be delocalized over multiple atoms, the conjugate base is more stable and the corresponding acid is stronger. For example, HNO₃ is a strong acid because nitrate (\(\text{NO}_3^-\)) is resonance-stabilized.
Inductive effects from electronegative substituents can significantly increase acid strength. In organic acids like halogen-substituted carboxylic acids, electronegative atoms pull electron density away from the O–H bond, increasing polarity and facilitating proton release. The effect decreases with increasing distance from the acidic site.
On the AP exam, comparing acid strengths based on periodic trends, electronegativity, resonance, and inductive effects is a common conceptual question. A solid grasp of these structural factors allows quick predictions about which acid in a set will be stronger and why.

Base Strength Trends

Base strength refers to the ability of a substance to accept protons or produce hydroxide ions in solution. For metal hydroxides of Group 1A and heavier Group 2A metals, strength is high because they dissociate completely in water to release \(\text{OH}^-\). These strong bases, such as NaOH, KOH, and Ba(OH)₂, are treated as fully ionized for pH and pOH calculations.
For binary compounds containing a metal and a less electronegative element (like alkali and alkaline earth metal oxides), basicity increases down a group as metallic character increases. Larger cations hold the hydroxide ion less tightly, allowing easier dissociation. This explains why CsOH is a stronger base than LiOH.
In molecular (non-metal) bases, basicity trends are opposite to acid strength trends. For example, in ammonia and amine derivatives, basicity increases with greater electron density on the nitrogen that is available to accept a proton. Electron-donating groups increase basicity, while electron-withdrawing groups decrease it.
Oxyanion bases show increased basicity when the central atom is less electronegative or has fewer oxygen atoms, because this reduces the pull on electron density from the basic oxygen site. This allows the oxygen to donate electrons more readily to a proton. For example, \(\text{PO}_4^{3-}\) is a stronger base than \(\text{HPO}_4^{2-}\).
On the AP exam, base strength questions often test recognition of periodic trends, the relationship between conjugate acid strength and base strength, and the effect of molecular structure on proton affinity. Identifying the strongest or weakest base in a set often requires applying both periodic and electronic structure concepts.

Acid-Base Reactions and Hydrolysis

Acid-base neutralization reactions occur when an acid reacts with a base to form water and a salt. For strong acid–strong base reactions, the resulting solution is neutral (pH ≈ 7 at 25°C) because neither ion hydrolyzes significantly. For strong acid–weak base or weak acid–strong base reactions, the conjugate of the weak species hydrolyzes, shifting the pH away from 7.
Salt hydrolysis occurs when the anion of a weak acid or the cation of a weak base reacts with water to produce \(\text{OH}^-\) or \(\text{H}^+\), respectively. For example, in NaF, \(\text{F}^-\) reacts with water to form HF and \(\text{OH}^-\), making the solution basic. In NH₄Cl, \(\text{NH}_4^+\) reacts with water to form NH₃ and \(\text{H}^+\), making the solution acidic.
Determining whether a salt solution is acidic, basic, or neutral requires examining both ions for hydrolysis potential. If both ions are from strong acids and bases, the solution is neutral. If one ion comes from a weak species, its hydrolysis will dominate the pH. If both ions are from weak species, the pH depends on the relative \(K_a\) and \(K_b\) values.
Amphiprotic ions, such as \(\text{HCO}_3^-\) and \(\text{HSO}_4^-\), can act as either acids or bases, so their pH effect depends on which process is stronger. This is determined by comparing \(K_a\) and \(K_b\) for the ion. The stronger tendency will determine whether the solution is acidic or basic.
On the AP exam, hydrolysis problems often require writing the relevant equilibrium reaction, identifying the species acting as the acid or base, and calculating pH using \(K_a\) or \(K_b\). Understanding hydrolysis is essential for predicting pH in solutions of salts and in buffer systems.

Properties of Buffers

Buffers are solutions that resist significant changes in pH when small amounts of acid or base are added. They work because they contain a weak acid and its conjugate base (or a weak base and its conjugate acid) in roughly comparable concentrations. When an acid is added, the conjugate base neutralizes it, and when a base is added, the weak acid neutralizes it.
The effectiveness of a buffer depends on both the absolute concentrations of the acid–base pair and their ratio. Higher total concentrations result in greater buffer capacity, meaning the solution can absorb more added acid or base without large pH changes. A 0.50 M buffer is more effective than a 0.05 M buffer of the same ratio.
Buffer pH is most resistant to change when the concentrations of the acid and conjugate base are equal, which occurs when pH ≈ pKa. At this point, the buffer can neutralize equal amounts of added strong acid or base with minimal change in pH, a condition called optimal buffering.
Buffers have limits—once the added strong acid or base exceeds the neutralizing capacity of either the acid or conjugate base component, the pH will change rapidly. This limit is referred to as the buffer capacity, and it is exceeded when the ratio of acid to base becomes too skewed.
In biological and environmental systems, buffers are crucial for maintaining stable pH levels. For example, the bicarbonate buffer system in human blood maintains pH near 7.4 by balancing carbonic acid (H₂CO₃) and bicarbonate ions (HCO₃⁻), preventing dangerous fluctuations from metabolic processes.

The Henderson–Hasselbalch Equation

The Henderson–Hasselbalch equation provides a convenient way to calculate the pH of a buffer solution without solving an ICE table: \[ \text{pH} = \text{p}K_a + \log\left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \] where \([\text{A}^-]\) is the concentration of conjugate base and \([\text{HA}]\) is the concentration of weak acid.
This equation shows that when the conjugate base and acid concentrations are equal, the log term becomes zero and pH = pKa. This is the point of maximum buffer efficiency, as small additions of acid or base cause minimal pH change. This is why buffer preparation often targets a ratio close to 1:1.
Changing the ratio of \([\text{A}^-]\) to \([\text{HA}]\) shifts the pH predictably. If \([\text{A}^-] > [\text{HA}]\), the pH will be above the pKa, and if \([\text{A}^-] < [\text{HA}]\), the pH will be below the pKa. This relationship is logarithmic, so doubling the base-to-acid ratio increases pH by about 0.3 units.
The Henderson–Hasselbalch equation is only valid when both the acid and base are present in significant amounts and the acid is weak (Ka < 1). It cannot be applied to strong acid–base systems because they fully dissociate and lack the equilibrium needed for buffering.
On the AP exam, this equation is frequently used in buffer design problems, where you may need to calculate the amounts of acid and conjugate base to mix for a desired pH. Understanding both the equation and the chemical reasoning behind it is key to applying it correctly.

pH and Solubility

The solubility of many ionic compounds is affected by the pH of the solution, especially if the compound contains an anion that can act as a weak base. When pH is lowered (solution becomes more acidic), additional H⁺ ions react with the basic anion, reducing its concentration and shifting the dissolution equilibrium toward more dissolution according to Le Châtelier’s Principle.
For example, calcium fluoride (\(\text{CaF}_2\)) dissolves to produce \(\text{F}^-\) ions, which can react with H⁺ to form HF. By removing \(\text{F}^-\) through protonation, the equilibrium shifts to dissolve more \(\text{CaF}_2\), increasing solubility. This effect is especially significant for salts of weak acids such as carbonates, phosphates, and sulfides.
In contrast, if the anion of the salt is the conjugate base of a strong acid (e.g., Cl⁻ from HCl), pH changes have little or no effect on solubility. This is because strong acid conjugate bases do not hydrolyze significantly, so no acid–base equilibrium can be exploited to increase solubility.
Raising the pH (making the solution more basic) can decrease the solubility of salts with basic anions by increasing the common ion concentration. For example, adding OH⁻ to a solution containing \(\text{Mg(OH)}_2\) will shift the dissolution equilibrium toward the solid phase, decreasing solubility due to the common ion effect.
On the AP exam, pH and solubility problems often require combining \(K_{sp}\) concepts with acid–base equilibria, sometimes using multiple equilibrium steps to predict how pH manipulation will change the solubility of a sparingly soluble salt. These questions test your ability to link solubility rules, equilibrium shifts, and acid–base chemistry in a single explanation.

Titrations and pH Curves

Titrations are analytical techniques used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration. The process involves gradually adding the titrant (known solution) to the analyte (unknown solution) until the reaction reaches the equivalence point, where the moles of acid and base are stoichiometrically equal. The pH changes during the titration can be tracked using a pH meter or an appropriate indicator.
The shape of the pH curve depends on the strengths of the acid and base involved. Strong acid–strong base titrations produce a steep pH change around the equivalence point, typically at pH ≈ 7. Weak acid–strong base titrations produce an equivalence point at pH > 7 due to the formation of a basic conjugate base, while weak base–strong acid titrations have equivalence points at pH < 7 due to the formation of an acidic conjugate acid.
During weak acid–strong base titrations, a buffer region appears before the equivalence point where the solution contains significant amounts of both the weak acid and its conjugate base. In this region, the pH can be calculated using the Henderson–Hasselbalch equation, and pH changes slowly despite titrant addition. This is a key concept for recognizing and analyzing buffer behavior in titration curves.
The equivalence point is distinct from the endpoint, which is the point where the indicator changes color. An ideal indicator has a color change range that matches the steep pH change of the titration curve around the equivalence point, ensuring accurate determination of the titration's completion.
On the AP exam, titration problems often test your ability to interpret pH curves, calculate pH at different stages (before, at, and after equivalence), and select appropriate indicators. You should be comfortable with both the qualitative shape of the curve and the quantitative calculations at specific points.

Indicators

Indicators are weak acids or bases whose conjugate forms have different colors, allowing them to signal pH changes visually. The color change occurs over a narrow pH range, typically around the indicator’s pKa ± 1, where both acid and base forms are present in appreciable amounts. This range is called the transition range.
The choice of indicator depends on the expected equivalence point pH of the titration. For example, phenolphthalein (transition range pH 8.3–10) is ideal for weak acid–strong base titrations, while methyl orange (transition range pH 3.1–4.4) is better for strong acid–weak base titrations. Using the wrong indicator can cause large measurement errors.
Indicators work because of structural changes in their molecules that alter light absorption. In the acidic form, certain bonds or conjugated systems dominate, giving one color, while in the basic form, the molecular structure shifts, altering electron distribution and resulting in a different color.
For strong acid–strong base titrations, many indicators can be used because the pH change at equivalence is extremely sharp, spanning several pH units. This flexibility is why phenolphthalein and bromothymol blue are both commonly used in these titrations.
On the AP exam, indicator questions may involve matching an indicator to a given pH curve, predicting the observed color change, or explaining why a particular indicator would or would not be suitable for a specific titration type.

Common Misconceptions

Many students believe that the strength of an acid or base is the same as its concentration, but strength refers to the degree of ionization, not the molarity. For example, a 0.10 M solution of acetic acid is weaker than a 0.10 M solution of hydrochloric acid because acetic acid only partially ionizes, while HCl completely ionizes. The AP exam often tests this distinction through conceptual or calculation-based questions involving \(K_a\) and percent ionization.
Another common error is thinking that the pH at the equivalence point of any titration is always 7. In reality, the pH at equivalence depends on the nature of the acid and base: strong acid–strong base titrations have equivalence pH ≈ 7, but weak acid–strong base titrations yield pH > 7, and weak base–strong acid titrations yield pH < 7. Misunderstanding this leads to incorrect indicator selection and pH predictions.
Some students incorrectly assume that adding more buffer components always increases buffering capacity without limit. While increasing concentrations of the conjugate acid-base pair can improve capacity, the buffer is still most effective when pH ≈ pKa. A highly concentrated buffer far from its pKa will not resist pH change as effectively as one centered near its pKa.
There is a misconception that pH changes linearly with acid or base addition, but pH is a logarithmic scale, so small changes in [H⁺] can cause large pH shifts. This is especially important when interpreting titration curves — the steep regions reflect exponential changes in hydrogen ion concentration, not uniform shifts.
In pH and solubility problems, students often overlook that only salts with basic anions show increased solubility in acidic solutions. Attempting to apply this logic to salts like NaCl leads to incorrect predictions because Cl⁻ is the conjugate base of a strong acid and does not react with H⁺. AP free-response questions frequently test this nuance by mixing solubility and acid-base equilibria.