Now that we've got precipitation reactions down, let's move on to the next primary type of chemical reactions! The second main type of reaction that you learn in this unit is the acid-base neutralization reaction.
Defining Acids & Bases
There are many different ways to define the behavior of acids and bases. The AP Chemistry curriculum focuses on the Brønsted-Lowry definition. When you think "Brønsted-Lowry," you should immediately focus on the transfer of a proton or hydrogen ion.
A Proton
First things first, how is a proton equivalent to a hydrogen ion?
A proton is a subatomic particle that is found in the nucleus of an atom. It has a positive electric charge and is one of the fundamental building blocks of matter.
Remember, an ion is an atom or molecule that has gained or lost one or more electrons, resulting in a net electric charge. A hydrogen ion is a hydrogen atom that has been stripped of its electron, resulting in a net positive charge: H+. Because a proton is identical to a hydrogen ion, the terms "proton" and "hydrogen ion" are often used interchangeably in chemistry.
Sometimes, you may even see H₃O⁺ in place of H⁺.
Brønsted-Lowry Definitions
Focusing on the transfer of a proton, acids are proton donors while bases are proton acceptors. What you basically see happen is a hydrogen ion being transferred from a substance, denoted as the acid, to another substance, denoted as the base.
Since acid-base reactions are just transfers of hydrogen ions, who says they can't go both ways? They usually can, although they do often shift in a certain direction. This begins to cover content that is discussed later in this course, so let's keep it simple!
If acid-base reactions can go back and forth, there must be an acid and a base on both the reactant and product side. This creates conjugate acid-base pairs. Looking at a chemical equation, you should be able to tell what the acid-base pairs are and pick out the conjugates.
Let's focus on the following example: H₂O + H₂S → H₃O⁺ + HS⁻.
First things first, what are the acid-base pairs?
First pair: H₂O and H₃O⁺
Second pair: H₂S and HS⁻
Now, which are the acids, and which are the bases? A quick way to know would be to figure out which compound in the pair has an additional hydrogen. Since H₃O⁺ has one more hydrogen than H₂O, it is the conjugate acid. This makes H₂O the base.
Try the second pair on your own! 😊
Comparing Strengths of Conjugate Acid-Base Pairs
An important principle in acid-base chemistry is the inverse relationship between the strengths of conjugate acid-base pairs:
- Strong acids have weak conjugate bases: When an acid readily donates its proton, the resulting conjugate base has little tendency to accept a proton back.
- Example: HCl (strong acid) → Cl⁻ (very weak conjugate base)
- Weak acids have stronger conjugate bases: When an acid holds onto its proton more tightly, the resulting conjugate base has a greater tendency to accept protons.
- Example: CH₃COOH (weak acid) → CH₃COO⁻ (stronger conjugate base)
- Strong bases have weak conjugate acids: When a base readily accepts protons, the resulting conjugate acid has little tendency to donate the proton back.
- Example: OH⁻ (strong base) → H₂O (very weak conjugate acid)
This inverse relationship helps predict reaction directions: the stronger acid will donate protons to the stronger base, forming the weaker acid and weaker base. The equilibrium favors the side with the weaker acid and base.
👉 If you'd like more practice with this concept and to learn more about the Brønsted-Lowry definitions, make sure to review our study guide about titrations.
Amphiprotic Substances
There are also these special agents called amphiprotic substances. They can both donate and accept protons! The most important example is H₂O, but NH₃ and HCO₃⁻ are also amphiprotic substances.
The reason why they are amphiprotic is that they have both a lone pair that can accept and bond with a proton and a transferable proton that they can donate as well.
Water's Unique Role in Acid-Base Reactions
Water plays an incredibly important role in many acid-base reactions, as its molecular structure allows it to accept protons from and donate protons to dissolved species. This dual capability makes water the universal medium for acid-base chemistry in aqueous solutions.
When an acid dissolves in water:
- Water molecules accept protons from the acid, forming H₃O⁺ (hydronium ions)
- Example: HCl + H₂O → H₃O⁺ + Cl⁻
When a base dissolves in water:
- Water molecules donate protons to the base, forming OH⁻ (hydroxide ions)
- Example: NH₃ + H₂O → NH₄⁺ + OH⁻
This is why we often see H₃O⁺ instead of just H⁺ in aqueous solutions - the protons don't exist freely but are always associated with water molecules. Water's ability to both donate and accept protons is essential for facilitating proton transfer reactions in solution.
Acid-Base Neutralization
A neutralization reaction occurs when an acid and base react to often form an ionic salt and liquid water. The H⁺ ion from the acid combines with the OH⁻ from the base to form H₂O (l). The basic form of the reaction is acid + base → salt + water.
You usually have to write out the chemical reaction. Let's say the two reactants are HNO₃ (aq) and KOH (aq), what are the products?
To make things easier on yourself, automatically write out H₂O (l) since you know that a proton is transferred to form water. Then, just combine the remaining ions, which would form the salt: HNO₃ (aq) + KOH (aq) → H₂O (l) + KNO₃ (?)
Soluble Salt?
Using solubility rules, is KNO₃ soluble? Or is it a precipitate? Any compound with NO₃ is soluble, so KNO₃ is in the aqueous state: HNO₃ (aq) + KOH (aq) → H₂O (l) + KNO₃ (aq)
Net Ionic Equation
We're back to net ionic equations! Remember, a net ionic equation is a chemical equation that shows only the species precipitating in a chemical reaction, omitting the spectator ions. For review on this subject, be sure to check out an earlier study guide in this unit that focuses primarily on net ionic equations.
⚠️ So far, we've been practicing writing net ionic equations for precipitation reactions. In doing so, we only dissociated soluble salts. Here, in neutralization reactions, we have to be really careful not to dissociate weak acids and bases. This is because they only partially dissociate into their constituent ions. Make sure you remember the strong acids and bases!
| Strong Acids | Strong Bases |
|---|---|
| HCl | CaOH |
| HBr | SrOH |
| HI | BaOH |
| HNO₃ | Group 1 metal + OH⁻ |
| H₂SO₄ | |
| HClO₃ | |
| HClO₄ |
Luckily, HNO₃ is a strong acid and KOH is a strong base, so we can dissociate both completely in the chemical equation. This makes the complete ionic equation the following: H⁺ (aq) + NO₃⁻ (aq) + K⁺ (aq) + OH⁻ (aq) → H₂O (l) + K⁺ (aq) + NO₃⁻ (aq)
We're so close to done! Now, all you have to do is eliminate the spectator ions (K⁺ and NO₃⁻), and you get the corresponding net ionic equation: H⁺ (aq) + OH⁻ (aq) → H₂O (l)
Concentration of Ions Question
With an acid-base neutralization reaction, you could also find the concentration of the ions. With this chemical reaction, you specifically focus on the concentrations of the hydrogen ion and hydroxide ion. In other words, what is [H⁺]? [OH⁻]?
Let's say we are given the following information and are expected to solve for the concentration of the hydrogen and hydroxide ions: 0.250 M and 28.0 mL of HNO₃ and 0.320 M and 53.0 mL of KOH.
Mole Calculations
Let's find the number of moles of HNO₃ and KOH using the equation for molarity.
Molarity = moles / volume in L - We have to convert the volumes we have into L by dividing by 1000.
HNO₃: 0.250 M = x moles / 0.0280 L
x = 0.00700 moles of HNO₃
KOH: 0.320 M = x moles / 0.0530 L
x = 0.0170 moles of KOH
Limiting Reactant
The limiting reactant is the reactant that there is less of. In this case, since there is a one-to-one ratio for all compounds, we can quickly identify HNO₃ as the LR.
Ion with a Concentration of Zero?
Since H⁺ is in both the limiting reactant and H₂O, it has a concentration of 0. The spectator ions cannot have a concentration of 0. Remember, they just help the reaction take place.
Half the question is done🥳! [H⁺] = 0.
[OH⁻]?
Finding the concentration of the excess compound is often the hardest part of the problem. First, we have to find the number of moles that reacted by converting the LR into the product. Again, since everything is one-to-one, we don't have to do extra stoichiometry. The number of moles that reacted is 0.00700.
Then, we simply subtract from the number of excess moles we started with, which is 0.0170 moles of KOH.
0.0170 - 0.00700 = 0.010 moles unreacted.
Last but not least, we need a volume! 28.0mL + 53.00mL = 0.081 L
0.010 moles unreacted / 0.081 L = 0.12 M of OH⁻
Final Answers
[H⁺] = 0
[OH⁻] = 0.12
Practice, practice, practice! It's just a lottttt of stoichiometry 🙃.
Net Ionic Equation Practice
Write the net ionic equation of a reaction between HNO₃ and Al(OH)₃.
Here are the steps you should take:
- Write out the products: H2O + Al(NO3)₃
- Balance the equation: 3HNO₃ + Al(OH)₃ → 3H2O + Al(NO₃)₃
- Write out the states of matter using solubility rules: 3HNO₃ (aq) + Al(OH)₃ (s) → 3H2O (l) + Al(NO₃)₃ (aq)1. Al(OH)₃ is insoluble! We cannot dissociate it in the next step
- Dissociate aqueous substances: 3H⁺ (aq) + 3NO₃ (aq) + Al(OH)₃ (s) → 3H₂O (l) + Al⁺³ (aq) + 3NO₃⁻ (aq)
- Identify spectator ions: 3H⁺ (aq) + 3NO₃ (aq) + Al(OH)₃ (s) → 3H₂O (l) + Al⁺³ (aq) + 3NO₃⁻ (aq)
- Cross out spectator ions: 3H⁺ (aq) + Al(OH)₃ (s) → 3H₂O (l) + Al⁺³ (aq)
Vocabulary
The following words are mentioned explicitly in the College Board Course and Exam Description for this topic.
| Term | Definition |
|---|---|
| aqueous solution | A solution in which water is the solvent. |
| Brønsted-Lowry acid | A species that donates a proton (H⁺) in a chemical reaction. |
| Brønsted-Lowry base | A species that accepts a proton (H⁺) in a chemical reaction. |
| conjugate acid-base pair | Two species that differ by one proton, where one is the acid form and the other is the base form of the same substance. |
| ionization | The process by which an acid or base separates into ions when dissolved in water. |
| proton transfer | The movement of a proton (H⁺) from one species to another in an acid-base reaction. |
Frequently Asked Questions
What is a Brønsted-Lowry acid and how is it different from other types of acids?
A Brønsted–Lowry acid is any species that donates a proton (H+); a Brønsted–Lowry base is any species that accepts a proton. In an acid–base reaction you always form a conjugate acid–base pair (acid → conjugate base after losing H+, base → conjugate acid after gaining H+). In aqueous chemistry this often shows up as H3O+ (hydronium) and OH−, and water is amphiprotic because it can donate or accept a proton. Strengths are compared with Ka and Kb; strong acids fully donate protons in water, weak acids only partially ionize. How it differs from other definitions: Arrhenius acids/bases are more limited—Arrhenius acids increase [H+] and bases increase [OH−] in water. Lewis acids/bases (electron-pair acceptors/donors) are broader but not tested on the AP Exam. For AP Chem focus on proton transfer, conjugate pairs, Ka/Kb, and aqueous reactions (see the Topic 4.8 study guide for a clear review: https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl). For extra practice try Fiveable’s AP Chem practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why do acids donate protons and bases accept them - what's actually happening?
By the Brønsted–Lowry definition, acids donate H+ and bases accept H+ because a proton transfer lowers the energy (raises stability) of one or both partners. Practically that means: an acid has an H attached to a polar bond (H—X) where the X is electronegative or the bond is weak, so the proton is relatively easy to pull off. A base has a lone pair or negative charge that can form a new bond to H+, so accepting a proton is favorable. After transfer you get a conjugate base and conjugate acid; the better stabilized the conjugate (by electronegativity, resonance, solvation by water), the stronger the original acid/base (Ka, Kb control equilibrium position). In water, H+ is actually H3O+ and water is amphiprotic—it can accept or donate protons (CED 4.8.A, keywords: hydronium, amphiprotic water). For more examples and AP-style practice, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and Unit 4 overview (https://library.fiveable.me/ap-chemistry/unit-4).
I'm confused about conjugate acid-base pairs - can someone explain this concept?
A conjugate acid-base pair is two species that differ by one proton (H+). When an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. Example: HCl (acid) → Cl− (conjugate base); NH3 (base) + H+ → NH4+ (conjugate acid). In water, remember H3O+ is the acid formed when something donates a proton to H2O, and OH− is the base when H2O donates a proton. Conjugate pairs tell you which direction proton transfer can go: a strong acid has a very weak conjugate base (HCl → Cl−), and a weak acid has a stronger conjugate base (acetic acid CH3COOH → CH3COO−). You’ll be asked to identify these on the AP exam under Brønsted–Lowry definitions and compare relative strengths using Ka or Kb (CED 4.8.A and 4.8.A.3). For extra practice, check the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and the 1000+ practice questions (https://library.fiveable.me/practice/ap-chemistry).
What's the difference between an acid and its conjugate base?
An acid and its conjugate base differ by one proton (H+). In Brønsted–Lowry terms an acid is a proton donor; after it donates H+ it becomes its conjugate base. Example: HCl (acid) → H+ + Cl−; Cl− is HCl’s conjugate base. Conversely, a base that accepts a proton becomes its conjugate acid (NH3 + H+ → NH4+). Conjugate pairs are linked: the stronger an acid (large Ka, fully dissociating strong acids), the weaker its conjugate base—a strong acid’s conjugate base has negligible tendency to reaccept H+. For weak acids (smaller Ka) the conjugate base is a stronger base (larger Kb) and can reaccept H+. In aqueous reactions remember water can act as both acid and base (amphiprotic), producing H3O+ and OH−. This conjugate-pair idea is tested on the AP exam when you identify acids/bases and compare strengths (see Topic 4.8 study guide on Fiveable: https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl). For extra practice try problems at Fiveable’s practice page (https://library.fiveable.me/practice/ap-chemistry).
How do I identify which species is the acid and which is the base in a reaction?
Use the Brønsted–Lowry definition: an acid donates a proton (H+), a base accepts a proton. To decide in any equation, ask “which species loses H+?” That one is the acid; “which gains H+?” is the base. Identify the conjugate acid–base pairs by comparing reactants and products: the acid becomes its conjugate base after losing H+, and the base becomes its conjugate acid after gaining H+. If the reaction is in water, remember H2O can act as either acid or base (amphiprotic). Example: HCl + H2O → H3O+ + Cl−. HCl donates H+ (acid), H2O accepts H+ (base); products H3O+ (conjugate acid) and Cl− (conjugate base). Strong acids fully donate H+ (their conjugate bases are weak). AP focus: only Brønsted–Lowry in aqueous solutions and identifying conjugate pairs (CED 4.8.A). For more practice and clear examples, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and extra problems (https://library.fiveable.me/practice/ap-chemistry).
Why does water act as both an acid and a base in some reactions?
Water can act as both an acid and a base because of its ability to donate or accept a proton (H+), which is exactly the Brønsted–Lowry idea you study in Topic 4.8. As a base, H2O accepts a proton to form H3O+ (hydronium): H2O + HCl → H3O+ + Cl−. As an acid, H2O donates a proton to form OH− (hydroxide): H2O + NH3 → OH− + NH4+. That dual role comes from water’s molecular structure and the fact we work in aqueous solutions—water is amphiprotic and participates in autoionization (2 H2O ⇌ H3O+ + OH−). On the AP exam you should identify species by proton transfer and name conjugate acid–base pairs (e.g., H2O/OH− or H3O+/H2O). For a quick topical review, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and more practice problems at (https://library.fiveable.me/practice/ap-chemistry).
What happens when HCl dissolves in water and how do I write the equation?
When HCl dissolves in water, it acts as a strong Brønsted–Lowry acid and donates a proton to water. The proton transfer is essentially complete, so write the reaction showing water as the proton acceptor: HCl(aq) + H2O(l) → H3O+(aq) + Cl−(aq) You can also show the net ionic/proton-transfer view as: H+(aq) + H2O(l) → H3O+(aq) Key points: HCl is the acid (proton donor), H2O is the base (proton acceptor). The conjugate acid–base pairs are HCl/Cl− and H3O+/H2O. Because HCl is a strong acid in water, it essentially fully ionizes, so the solution contains H3O+ and Cl−. This is exactly the Brønsted–Lowry proton-transfer idea tested in Topic 4.8 (use H3O+ rather than just H+ on the AP exam when showing water’s role). For a quick refresher, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and practice problems (https://library.fiveable.me/practice/ap-chemistry).
How do I know if an acid or base is strong or weak just by looking at it?
Short answer: you usually tell strong vs. weak by memorized families, solubility rules, and whether the acid/base fully ionizes in water (Brønsted–Lowry idea—proton donor/acceptor). Quick rules you should memorize for AP: - Strong acids (completely ionize): HCl, HBr, HI, HNO3, HClO4, and the first proton of H2SO4. Anything else is treated as a weak acid unless told otherwise. - Strong bases (strongly ionize/produce OH−): soluble Group 1 hydroxides (LiOH, NaOH, KOH, …) and the heavier Group 2 hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2). Most other bases (like NH3, organic amines) are weak. - If you’re unsure, think: does it fully dissociate in water? If yes → strong; if partial → weak. On the exam you may also be given Ka or Kb (large Ka/Kb = strong; small = weak). For more AP-aligned review and examples, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Can you explain proton transfer reactions step by step?
A proton-transfer (Brønsted–Lowry) reaction is just a stepwise handoff of H+ from a donor (acid) to an acceptor (base). Step 1: identify the acid (proton donor) and the base (proton acceptor). Step 2: show the acid losing H+ and the base gaining H+. Example in water: HCl + H2O → H3O+ + Cl−. Here HCl is the acid, H2O is the base, H3O+ is the conjugate acid, and Cl− is the conjugate base. Step 3: check conjugate pairs—acid ⇄ conjugate base, base ⇄ conjugate acid. Step 4: consider strength and equilibrium: strong acids (like HCl) ionize nearly completely so equilibrium lies far right; weak acids establish an equilibrium quantified by Ka (and Kb for bases). Step 5: for aqueous problems, write Ka or Kb and set up an ICE table if you need concentrations or pH. Use these steps for identifying species, conjugate pairs, and predicting direction—all directly tied to CED Topic 4.8 (see the study guide on Fiveable: https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl). For extra practice, try problems at Fiveable (https://library.fiveable.me/practice/ap-chemistry).
I don't understand how to identify conjugate acid-base pairs in equations - help?
Start by asking: who donates a proton (H+) and who accepts it? A Brønsted–Lowry acid donates H+; a base accepts H+. Conjugate acid–base pairs are the two species related by the gain or loss of ONE proton—they differ by H+. Quick steps: 1. Identify the proton transfer in the equation. 2. The species that loses H+ is the acid; its partner after losing H+ is its conjugate base. 3. The species that gains H+ is the base; its partner after gaining H+ is its conjugate acid. Examples: - HCl(aq) + H2O(l) → Cl−(aq) + H3O+(aq) Conjugate pairs: HCl / Cl− and H3O+ / H2O. - NH3(aq) + H2O(l) → NH4+(aq) + OH−(aq) Pairs: NH3 / NH4+ and H2O / OH−. Remember: only aqueous reactions are tested on the AP exam, and water can act as either acid or base (amphiprotic). For a clear walkthrough and practice, check the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and grab more practice problems at (https://library.fiveable.me/practice/ap-chemistry).
Why is water so important in acid-base reactions compared to other solvents?
Water matters in acid–base chemistry because AP Chem focuses on reactions in aqueous solution and water is amphiprotic—it can both donate and accept protons. That means dissolved acids donate H+ to water to form H3O+ (hydronium) and bases accept H+ from water to form OH−. Water’s autoionization (Kw ≈ 1.0×10−14 at 25°C) sets the background concentrations of H3O+ and OH−, which you use to compare acid and base strengths (Ka, Kb) and identify conjugate acid–base pairs. In other solvents you don’t get the same universal proton-transfer partner or the same predictable equilibrium (so many AP tasks stick to aqueous cases). For more practice and clear examples tied to the CED keywords (hydronium, hydroxide, amphiprotic water, autoionization), check the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl). For broader unit review and thousands of practice problems, see the Unit 4 page (https://library.fiveable.me/ap-chemistry/unit-4) and practice bank (https://library.fiveable.me/practice/ap-chemistry).
What does it mean when we say water can accept and donate protons?
Saying water can both accept and donate protons means it’s amphiprotic under the Brønsted–Lowry definition: an acid is a proton (H+) donor and a base is a proton acceptor. Water donates a proton when it acts as an acid: H2O → H+ + OH− (more realistically H2O → H3O+? No—when donating it forms OH−). It accepts a proton when it acts as a base: H2O + H+ → H3O+. Those two reactions show conjugate acid–base pairs (H2O/OH− and H3O+/H2O). This behavior underlies water’s autoionization (2 H2O ⇌ H3O+ + OH−) and explains why H3O+ and OH− are the key species in aqueous acid–base problems on the AP: identify donors/acceptors and their conjugates. For a clear AP-aligned review, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and more practice at (https://library.fiveable.me/practice/ap-chemistry).
How do I compare the relative strengths of conjugate acid-base pairs?
Compare conjugate pairs by asking: which member more readily donates or accepts a proton in water? Key rules (aqueous, Brønsted–Lowry): - Stronger acid → larger Ka (smaller pKa) → its conjugate base is weaker. Example: HCl is a strong acid (Ka huge), so Cl− is a very weak base. - Stronger base → larger Kb (smaller pKb) → its conjugate acid is weaker. Example: OH− is a very strong base; H2O is its conjugate acid (weak). - Use Ka and Kb: for a conjugate pair, Ka · Kb = Kw (1.0×10−14 at 25°C). So if you know Ka for an acid, compute Kb for its conjugate base. - Compare pKa values: lower pKa = stronger acid; its conjugate base is correspondingly weaker. - Predict equilibrium: proton transfer favors formation of the weaker acid and weaker base. For practice and AP-aligned examples, see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl), the Unit 4 overview (https://library.fiveable.me/ap-chemistry/unit-4), and extra practice problems (https://library.fiveable.me/practice/ap-chemistry).
What's the pattern for writing acid-base reactions with water?
Use Brønsted–Lowry logic: acids donate H+, bases accept H+. In water that means follow two simple patterns: - Acid + H2O → conjugate base + H3O+ - Example: HCl(aq) + H2O(l) → H3O+(aq) + Cl−(aq) - Base + H2O → conjugate acid + OH− - Example: NH3(aq) + H2O(l) → NH4+(aq) + OH−(aq) - Also remember autoionization: 2 H2O ⇌ H3O+ + OH− When you write reactions, always show water acting as the proton donor or acceptor (amphiprotic). Identify conjugate acid–base pairs (HCl/Cl−, NH4+/NH3). For strong acids/bases, write them fully ionized in water; for weak ones, show the equilibrium arrow and use Ka or Kb. This is exactly what the AP CED emphasizes (Brønsted–Lowry, hydronium/hydroxide, conjugate pairs). For a quick refresher see the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl) and practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why don't we need to know Lewis acids for the AP exam and what's the difference?
Short answer: The AP course and exam only test Brønsted–Lowry acid–base ideas (proton donors and acceptors) because Topic 4.8 and the CED focus on proton transfer in aqueous solutions—identify acids/bases, conjugate pairs, H3O+ and OH−, Ka/Kb, strong vs weak, etc. The CED explicitly says Lewis acid–base concepts will not be assessed on the AP Exam. Difference in one line: Brønsted–Lowry = acid donates H+ and base accepts H+ (works in water, conjugate pairs). Lewis = acid accepts an electron pair and base donates an electron pair (broader—covers reactions that don’t involve protons, coordination chemistry, etc.). Lewis is useful in general chemistry but not required for AP Chem exam questions. If you want a quick review of the AP scope for this topic, check the Topic 4.8 study guide (https://library.fiveable.me/ap-chemistry/unit-4/intro-acid-base-reactions/study-guide/idvZ7Ve4pFo8gFIyYfMl). For more practice on what AP actually tests, use the unit page (https://library.fiveable.me/ap-chemistry/unit-4) and the 1000+ practice problems (https://library.fiveable.me/practice/ap-chemistry).