Acids, Bases, and pH Study Pack
Kibin's free study pack on Acids, Bases, and pH includes a 6-section study guide, 25 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Acids, Bases, and pH Study Guide
Master the Brønsted-Lowry theory of proton donors and acceptors, conjugate acid-base pairs, and the amphoteric nature of water. From autoionization and pH calculations to Ka, Kb, and strong vs. weak dissociation, this pack covers the core mechanics of acid-base chemistry.
Key Takeaways
- •Brønsted-Lowry theory defines acids as proton (H⁺) donors and bases as proton acceptors, allowing acid-base behavior to be described in any solvent, not just water.
- •Every Brønsted-Lowry acid-base reaction produces a conjugate base (from the acid) and a conjugate acid (from the base), forming conjugate acid-base pairs.
- •Water is amphoteric — it can donate or accept protons depending on its reaction partner, acting as either an acid or a base.
- •The autoionization of water produces equal concentrations of H₃O⁺ and OH⁻ at 25 °C, each at 1.0 × 10⁻⁷ M, giving pure water a pH of exactly 7.
- •pH is defined as the negative base-10 logarithm of the hydronium ion concentration: pH = −log[H₃O⁺]; values below 7 indicate acidic solutions and values above 7 indicate basic solutions.
- •Strong acids and strong bases dissociate completely in water, while weak acids and weak bases establish equilibria described by Ka and Kb, respectively.
- •The acid dissociation constant Ka measures acid strength quantitatively — a higher Ka means a stronger acid that more readily donates protons.
Defining Acids and Bases: The Brønsted-Lowry Model
The Brønsted-Lowry model, proposed independently by Johannes Brønsted and Thomas Lowry in 1923, provides the most widely used framework for understanding acid-base chemistry at the molecular level.
Core Definitions in Brønsted-Lowry Theory
- •A Brønsted-Lowry acid is any species that donates a proton (H⁺) to another species during a chemical reaction.
- •A Brønsted-Lowry base is any species that accepts a proton from an acid.
- •Because a proton cannot exist freely in solution, every acid donation requires a base to receive it — acid-base reactions always occur in matched pairs.
Contrast with the Arrhenius Model
- •The earlier Arrhenius model restricted acids to substances that release H⁺ in water and bases to substances that release OH⁻ in water.
- •Brønsted-Lowry expands this scope: ammonia (NH₃) qualifies as a base because it accepts protons, even though it contains no OH⁻.
- •This broader definition allows acid-base chemistry to be analyzed in non-aqueous solvents and explains reactions the Arrhenius model cannot.
Conjugate Acid-Base Pairs and Proton Transfer
When an acid donates a proton, the species it becomes is called its conjugate base; when a base accepts a proton, the species it becomes is called its conjugate acid — these paired relationships are central to tracking proton movement in any reaction.
Formation of Conjugate Pairs
- •After HCl donates H⁺ to water, HCl becomes Cl⁻ — the conjugate base of HCl.
- •Water accepts that H⁺ to become H₃O⁺ (hydronium ion) — the conjugate acid of water.
- •Every Brønsted-Lowry reaction contains exactly two conjugate pairs: the original acid/its conjugate base, and the original base/its conjugate acid.
Relative Strength Within a Conjugate Pair
- •A strong acid (one that donates protons readily) always has a weak conjugate base (one that holds protons tightly and rarely donates them back).
- •Conversely, a weak acid has a relatively strong conjugate base.
- •This inverse relationship means that if you know the strength of an acid, you can immediately infer the strength of its conjugate base.
Amphoteric Species
- •Water is the classic amphoteric species — it acts as a base when reacting with HCl (accepting H⁺ to form H₃O⁺) and as an acid when reacting with NH₃ (donating H⁺ to form OH⁻ and NH₄⁺).
- •The bicarbonate ion (HCO₃⁻) is another common amphoteric species, capable of donating or accepting a proton depending on its reaction partner.
Water Autoionization and the Ion-Product Constant
Even in pure water, a small fraction of molecules transfer protons to one another in a process called autoionization, which establishes the baseline relationship between H₃O⁺ and OH⁻ concentrations that underlies all aqueous acid-base chemistry.
The Autoionization Equilibrium
- •Two water molecules can react so that one acts as the acid (donating H⁺) and the other acts as the base (accepting H⁺), producing one H₃O⁺ ion and one OH⁻ ion.
- •This equilibrium is written: 2 H₂O ⇌ H₃O⁺ + OH⁻.
- •At 25 °C, the equilibrium lies far to the left; only about 1 in 10 million water molecules is ionized at any moment.
The Ion-Product Constant for Water (Kw)
- •The equilibrium expression for autoionization gives Kw = [H₃O⁺][OH⁻].
- •At 25 °C, Kw = 1.0 × 10⁻¹⁴; in pure water, both [H₃O⁺] and [OH⁻] equal 1.0 × 10⁻⁷ M.
- •Kw holds for any dilute aqueous solution at 25 °C — if acid is added and [H₃O⁺] rises above 1.0 × 10⁻⁷ M, [OH⁻] must fall proportionally to keep the product equal to 1.0 × 10⁻¹⁴.
- •Kw increases with temperature, meaning hot water is slightly more ionized than cold water.
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What is the core definition of a Brønsted-Lowry acid?
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Brønsted-Lowry Acid-Base Theory
Explain the Brønsted-Lowry definitions of acids and bases in your own words. How does this model differ from the earlier Arrhenius model, and why is the Brønsted-Lowry approach considered more useful?
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