Deprotonation is the loss of a proton (H⁺) from a molecule or ion, converting a Brønsted-Lowry acid into its conjugate base. In AP Chem Unit 8, it's the fundamental event in every acid-base reaction, including water's autoionization (Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C).
Deprotonation is what happens when an acid does its job. A Brønsted-Lowry acid donates a proton (H⁺) to a base, and the species left behind is the conjugate base. Take away one H and one unit of positive charge, and you've deprotonated. HCl becomes Cl⁻. HC₂H₃O₂ becomes C₂H₃O₂⁻. NH₄⁺ becomes NH₃.
The proton never just floats off on its own in solution. Something has to grab it, and in water that something is usually a water molecule, which becomes hydronium ion (H₃O⁺). That's why deprotonation and protonation are two halves of one event called proton transfer. Even pure water does this to itself. In autoionization, one water molecule deprotonates another, producing H₃O⁺ and OH⁻ with the equilibrium constant Kw. The AP Exam prefers you write H₃O⁺(aq), though H⁺(aq) is also accepted.
Deprotonation lives in Topic 8.1 (Introduction to Acids and Bases) and underpins learning objective 8.1.A, which asks you to calculate pH and pOH using Kw and the concentrations of species in water. The autoionization of water is literally water deprotonating water, and the result is the equation Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C, which gives you pH + pOH = 14 at that temperature.
But the payoff goes way beyond 8.1. Every weak-acid equilibrium (Ka), every buffer, and every titration curve in Unit 8 is just deprotonation happening to a greater or lesser extent. Strong acids deprotonate completely; weak acids deprotonate only partially, and Ka measures how far that goes. If you can spot which species lost an H⁺ and which gained one, you can set up almost any Unit 8 problem.
Keep studying AP® Chemistry Unit 8
Proton Transfer and Brønsted-Lowry Theory (Unit 8)
Deprotonation is one side of a proton transfer. The acid deprotonates while the base protonates, in the same instant. The Brønsted-Lowry definitions are built entirely on this single move, so identifying who lost the H⁺ tells you who the acid was.
Conjugate Base (Unit 8)
The conjugate base is simply the leftover after deprotonation. Acid minus H⁺ equals conjugate base, which means the pair differs by exactly one proton. HF and F⁻ are a conjugate pair; HF and F₂ are not, and the exam loves testing that distinction.
Kw and the Autoionization of Water (Unit 8)
Water deprotonates itself. One H₂O hands a proton to another, making H₃O⁺ and OH⁻. That equilibrium gives you Kw = 1.0 × 10⁻¹⁴ at 25°C and the relationship pH + pOH = 14, the backbone of LO 8.1.A calculations.
Acid-Base Reactions and Net Ionic Equations (Unit 4)
You first met deprotonation back in Unit 4 when classifying reaction types. Writing the net ionic equation for an acid-base reaction means showing exactly which species loses H⁺ and which species picks it up.
You won't usually see a question that says "define deprotonation." Instead, the exam tests whether you can use it. Multiple-choice stems ask you to identify the conjugate base formed when an acid donates a proton, or to label the acid, base, conjugate acid, and conjugate base in an equation. FRQs ask you to write equations for acids reacting with water, where the whole answer hinges on showing the proton leaving the acid and landing on H₂O to form H₃O⁺. No released FRQ uses the word "deprotonation" verbatim, but the skill shows up constantly in weak-acid equilibrium setups, buffer questions, and particulate diagrams where you must draw which species have lost their H⁺. One concrete tip is that the exam accepts H⁺(aq) but prefers H₃O⁺(aq), so write the hydronium version when you can.
They're mirror images of the same event. Deprotonation is losing H⁺ (acid becomes conjugate base), while protonation is gaining H⁺ (base becomes conjugate acid). In any Brønsted-Lowry reaction both happen simultaneously, so a common exam trap is mixing up which species did which. Track the H: the formula that has one fewer H and one more negative charge after the reaction got deprotonated.
Deprotonation means a molecule or ion loses a proton (H⁺), turning a Brønsted-Lowry acid into its conjugate base.
A conjugate acid-base pair differs by exactly one H⁺, so if two formulas differ by anything more, they are not a conjugate pair.
In water, the released proton is picked up by H₂O to form hydronium ion (H₃O⁺), which is the preferred symbol on the AP Exam over H⁺(aq).
Water's autoionization is water deprotonating water, producing H₃O⁺ and OH⁻ with Kw = 1.0 × 10⁻¹⁴ at 25°C, which gives pH + pOH = 14.
Strong acids deprotonate completely in water, while weak acids deprotonate only partially, and Ka measures the extent of that deprotonation.
Deprotonation and protonation always happen together as one proton transfer, so every acid-base reaction has both.
Deprotonation is the loss of a proton (H⁺) from a molecule or ion. The species that loses the proton is acting as a Brønsted-Lowry acid, and what's left behind is its conjugate base, like HF becoming F⁻.
Not exactly. Dissociation is a general term for a compound splitting into ions (like NaCl separating in water), while deprotonation specifically means losing an H⁺ to a base. Acid "dissociation" in water is really a proton transfer where H₂O accepts the proton and becomes H₃O⁺.
Deprotonation is losing H⁺ and protonation is gaining H⁺. In any Brønsted-Lowry reaction they happen at the same time, since the acid's lost proton is exactly the proton the base gains.
No. Weak acids deprotonate too, just partially, and that partial extent is what Ka measures. Even pure water deprotonates itself through autoionization, which is why Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C.
Both are accepted, but the CED says hydronium ion and H₃O⁺(aq) are preferred. Writing H₃O⁺ also shows you understand that the proton transfers to a water molecule rather than floating free in solution.
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