16.1 Introduction to Acids and Bases
PRE-HEALTH PREP
Intro to Acids and Bases
All Three Definitions
General Chemistry 16.1
Lesson Overview
Acids and bases are classified by three progressively broader definitions: Arrhenius (limited to aqueous solutions), Brønsted-Lowry (proton donors and acceptors), and Lewis (electron pair acceptors and donors). Conjugate acid-base pairs form when an acid donates a proton to become its conjugate base, or a base accepts a proton to become its conjugate acid, with amphiprotic species like bicarbonate capable of acting as either. Strong acids and bases dissociate completely in water, while weak acids and bases—including carboxylic acids and ammonia—dissociate only partially, typically less than 5%.
Table of Contents
What Are the Three Definitions of Acids and Bases?
In general chemistry, acids and bases are defined three different ways: Arrhenius, Brønsted-Lowry, and Lewis. These three definitions arrived in roughly historical order, and each one is progressively more complete than the last. Not all chemists throughout history have agreed on what makes something an acid or a base, which is exactly why three separate frameworks exist.
The Arrhenius definition is the most restrictive — it only applies to aqueous solutions, so it misses certain acids and bases entirely. Brønsted-Lowry broadens the scope to include reactions in any solvent or even in the gas phase. Lewis then delivers the ultimate definition: every acid and base recognized in modern chemistry satisfies Lewis's criteria, though the other two definitions fall short.
| Definition | Acid | Base | Scope |
|---|---|---|---|
| Arrhenius | Increases H3O+ in water | Increases OH− in water | Aqueous solutions only |
| Brønsted-Lowry | H+ (proton) donor | H+ (proton) acceptor | Any solvent or gas phase |
| Lewis | Electron pair acceptor | Electron pair donor | All acid-base reactions |
Key Relationship Between the Three Definitions
Every Arrhenius acid is also a Brønsted-Lowry acid, and every Brønsted-Lowry acid is also a Lewis acid. The reverse is not always true. Lewis has the most complete definition — if something is an acid by any criterion, it is guaranteed to be a Lewis acid.
What Is the Arrhenius Definition of Acids and Bases?
Arrhenius kept things simple: everything had to be in water. An Arrhenius acid is any substance that increases the concentration of hydronium ions (H3O+) when dissolved in water. An Arrhenius base is any substance that increases the concentration of hydroxide ions (OH−) when dissolved in water.
Arrhenius Acids: Hydronium Formation
Consider the most famous acid of them all: HCl (hydrochloric acid). Found in chemistry labs, swimming pools (where it goes by muriatic acid), and even your own stomach, HCl is the quintessential acid. When placed in water, HCl donates an H+ ion to a water molecule, forming hydronium:
HCl + H2O → H3O+ + Cl−
Anything that produces hydronium when dissolved in water qualifies as an Arrhenius acid. You can also think of it as an H+ donor in water — but the critical requirement is that it must occur in water.
Arrhenius Bases: Hydroxide Formation
On the base side, NaOH (sodium hydroxide) is the most famous example. It simply dissociates in water to release hydroxide ions:
NaOH → Na+ + OH−
The H3O+ and OH− Balance
In any aqueous solution, H3O+ and OH− coexist. Pure water has equal concentrations of both. Adding an acid tips the balance toward more H3O+ than OH−. Adding a base tips it the other way. This interplay between hydronium and hydroxide is central to aqueous acid-base chemistry.
What Is the Brønsted-Lowry Definition of Acids and Bases?
Acids and Bases as Proton Donors and Acceptors
Brønsted and Lowry expanded the playing field. The Brønsted-Lowry definition does not require water — acid-base reactions can occur in any solvent or even in the gas phase. A Brønsted-Lowry acid is simply an H+ donor, and a Brønsted-Lowry base is an H+ acceptor.
Why do chemists call H+ a "proton"? A hydrogen atom consists of one proton and one electron. Strip away that electron to form H+, and all that remains is the proton itself. So "proton donor" and "H+ donor" mean exactly the same thing.
Applying Brønsted-Lowry to Aqueous Reactions
Revisiting the HCl-in-water reaction through a Brønsted-Lowry lens: HCl donates the H+ (making it the acid), and water accepts the H+ to become H3O+ (making water the base). This reaction satisfies both the Arrhenius and Brønsted-Lowry definitions simultaneously.
Example: HCl in Water (Brønsted-Lowry Perspective)
HCl + H2O → H3O+ + Cl−
Acid: HCl (donates H+)
Base: H2O (accepts H+)
Beyond Water: Gas-Phase Acid-Base Reactions
Here is where Brønsted-Lowry truly shines. HCl and ammonia (NH3) can react in the gas phase — no water involved at all:
HCl(g) + NH3(g) → NH4Cl(s)
HCl donates a proton to become Cl−, and NH3 accepts that proton to become NH4+. Arrhenius would never call this an acid-base reaction because there is no water. Brønsted-Lowry has no such limitation.
Example: HCl and Ammonia in the Gas Phase
Acid: HCl (proton donor — loses H+ to become Cl−)
Base: NH3 (proton acceptor — gains H+ to become NH4+)
The Most Useful Working Definition
Brønsted-Lowry is the workhorse definition for identifying acids and bases in general chemistry. Follow the proton: whoever gives away the H+ is the acid, and whoever receives it is the base. This approach is more intuitive than tracking electron pairs (Lewis), and it works in far more situations than requiring aqueous solutions (Arrhenius).
What Is the Lewis Definition of Acids and Bases?
Electron Pair Donors and Acceptors
Lewis flipped the perspective from protons to electrons. A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor. The reasoning: in every acid-base reaction a new bond forms, and that bond requires a pair of electrons. The species that supplies those electrons is the Lewis base; the species that accepts them is the Lewis acid.
When that new bond forms to an H+ ion, the reaction qualifies as both Lewis and Brønsted-Lowry. When the new bond forms to something other than hydrogen, only Lewis's definition applies.
Lewis Structures Reveal the Electron Transfer
To identify Lewis acids and bases, draw the Lewis structures and locate the new bond in the products. Then ask: where did the bonding electrons come from?
Example: NH3 + HCl Through Lewis's Lens
The new bond forms between nitrogen and hydrogen. Nitrogen supplies the bonding electrons from its lone pair, making NH3 the electron pair donor (Lewis base). HCl contributes nothing to the bond — it merely accepts the electrons, making it the electron pair acceptor (Lewis acid).
Notice this matches the Brønsted-Lowry classification: HCl is still the acid and NH3 is still the base. The definitions agree whenever a proton is being transferred.
Uniquely Lewis Acids: Boron Compounds and Metal Ions
Some species qualify as acids only under the Lewis definition — neither Arrhenius nor Brønsted-Lowry would classify them as acids. Two major categories stand out:
- Boron compounds — Boron famously lacks a filled octet, which means it has an empty orbital ready to accept an electron pair. Compounds like BF3 and BH3 are classic Lewis acids.
- Metal ions — Cations such as Al3+ and Fe3+ possess empty valence orbitals, making them excellent electron pair acceptors and therefore Lewis acids.
The Hallmark of a Lewis Acid: Empty Orbitals
If a species has an empty orbital available to accept an electron pair, it can act as a Lewis acid. For uniquely Lewis acids (those that don't satisfy the other definitions), think boron compounds and metal ions.
The Hallmark of a Lewis Base: Lone Pairs
To act as a Lewis base, a species needs one thing: a lone pair of electrons. That lone pair can be donated to form a new bond. If the lone pair bonds to an H+, the species is both a Lewis base and a Brønsted-Lowry base. If it bonds to any other atom (like boron), it is a Lewis base but not a Brønsted-Lowry base.
Example: NH3 + BF3 — A Uniquely Lewis Acid-Base Reaction
When ammonia reacts with boron trifluoride, nitrogen donates its lone pair to boron's empty orbital. The new bond forms between N and B, with nitrogen gaining a positive formal charge and boron gaining a negative formal charge. No proton is transferred, so this is exclusively a Lewis acid-base reaction.
Summary: How the Three Definitions Compare
You should know each definition verbatim, but you also need to apply them. Brønsted-Lowry will be your go-to for identifying acids and bases throughout most of this chapter. However, always remember that Lewis holds the ultimate definition — if a substance is an acid under any definition, it is guaranteed to be a Lewis acid. There is a 100% chance.
What Are Conjugate Acid-Base Pairs?
Defining Conjugate Acid-Base Pairs
When an acid donates a proton, it transforms into its conjugate base. When a base accepts a proton, it transforms into its conjugate acid. These paired species — linked by the gain or loss of a single H+ — are called a conjugate acid-base pair. They typically exist in equilibrium, and their relationship determines the strength of the acid or base.
Identifying conjugate pairs relies on the Brønsted-Lowry definition: follow the proton.
Finding the Conjugate Base of an Acid
To find the conjugate base of any acid, simply have it act as an acid: donate an H+. Whatever remains is the conjugate base.
Example: Conjugate Base of HCl
HCl donates H+. Remove one H and one positive charge from the formula:
HCl → H+ + Cl−
Cl− is the conjugate base of HCl. After HCl acts as an acid, it becomes its conjugate base.
Finding the Conjugate Acid of a Base
To find the conjugate acid of any base, have it act as a base: accept an H+. The resulting species is the conjugate acid.
Example: Conjugate Acid of NH3
NH3 accepts H+. Add one H and one positive charge to the formula:
NH3 + H+ → NH4+
NH4+ is the conjugate acid of ammonia.
Amphiprotic Species: Bicarbonate
Some species can act as either an acid or a base — they are amphiprotic. Think of "amphi" like amphibian: an animal at home on land or in water. An amphiprotic species is at home donating or accepting protons.
This creates a potential exam trap. An amphiprotic species has both a conjugate acid and a conjugate base, and you need to know which one is being asked for.
Example: Bicarbonate (HCO3−) Is Amphiprotic
Conjugate base of HCO3−: Have bicarbonate donate H+. Losing one H and one positive charge gives CO32− (carbonate ion).
Conjugate acid of HCO3−: Have bicarbonate accept H+. Gaining one H and one positive charge gives H2CO3 (carbonic acid).
Exam Trap: The Conjugate Base of Hydroxide
Hydroxide (OH−) is overwhelmingly associated with being a base, so students rarely consider it acting as an acid. But if forced to donate H+, OH− becomes O2− (the oxide ion). This requires an exceptionally strong base to pull off, and you may encounter it in organic chemistry. The conjugate acid of OH− is simply H2O — water.
Identifying Conjugate Pairs in a Reversible Reaction
Most acid-base reactions are reversible, so both the forward and reverse directions contain an acid-base pair. Consider the reaction between acetic acid and ammonia:
CH3COOH + NH3 ⇌ CH3COO− + NH4+
Forward direction: CH3COOH donates H+ (acid) and NH3 accepts it (base).
Reverse direction: NH4+ donates H+ (acid) and CH3COO− accepts it (base).
The two conjugate pairs are: CH3COOH / CH3COO− and NH4+ / NH3. In any Brønsted-Lowry acid-base reaction, you should be able to identify the acid and base on both the reactant side and the product side.
Practice: Conjugate Acid/Base Pairs
For each species below, identify its conjugate partner. Remember: to find a conjugate base, remove H+; to find a conjugate acid, add H+. Note that HCO3− appears in both columns because it is amphiprotic.
| CONJUGATE ACID | CONJUGATE BASE |
|---|---|
| HCl | |
| NH3 | |
| HCO3− | |
| HCO3− | |
| OH− | |
| OH− |
What Is the Difference Between Strong and Weak Acids and Bases?
Strong vs. Weak: It Comes Down to Dissociation
The difference between strong and weak comes down to dissociation. Strong acids and bases dissociate effectively 100% in water. Weak acids and bases dissociate only partially — typically less than 5%, and often far less (values like 0.01% are not uncommon).
The Seven Strong Acids
Memorize these seven strong acids. Any acid not on this list should be treated as a weak acid by default.
| Strong Acid | Name |
|---|---|
| HCl | Hydrochloric acid |
| HBr | Hydrobromic acid |
| HI | Hydroiodic acid |
| HNO3 | Nitric acid |
| H2SO4 | Sulfuric acid |
| HClO4 | Perchloric acid |
| HClO3 | Chloric acid* |
*HClO3 appears on some lists but not others — it sits right at the boundary of strong acid behavior.
The Leveling Effect
In water, all strong acids dissociate completely, so the only acid actually present in solution is H3O+. This means all strong acid solutions appear to have the same acid strength in water. To distinguish their true relative strengths, you would need to use a different solvent. This phenomenon is called the leveling effect.
The Strong Bases
Strong bases are metal hydroxides from Group 1 and most of Group 2:
| Group 1 Hydroxides | Group 2 Hydroxides |
|---|---|
| LiOH (lithium hydroxide) | Mg(OH)2 (magnesium hydroxide) |
| NaOH (sodium hydroxide) | Ca(OH)2 (calcium hydroxide) |
| KOH (potassium hydroxide) | Sr(OH)2 (strontium hydroxide) |
| RbOH (rubidium hydroxide) | Ba(OH)2 (barium hydroxide) |
| CsOH (cesium hydroxide) |
Group 2 Hydroxides: Strong but Not Very Soluble
The Group 2 hydroxides are indeed strong bases — whatever dissolves will dissociate completely. However, they are not very soluble. Barium hydroxide maxes out around 0.1 M, and calcium hydroxide around 0.01 M. The amount that dissolves dissociates fully, but you cannot dissolve very much in the first place.
H2SO4: A Special Case Among Strong Acids
Sulfuric acid is diprotic — it has two protons to donate. The first H+ dissociates 100%, making it a strong acid for that first proton. However, the second H+ is only weakly acidic and dissociates partially. This distinction becomes important in pH calculations later in the chapter.
HF Is a Weak Acid (but Extremely Dangerous)
A common misconception: students assume hydrofluoric acid (HF) must be strong because it is so dangerous. In reality, HF is a weak acid — it is not on the strong acids list. HI, HBr, and HCl are all strong acids, but HF is not.
Safety Note: HF Is Deadly
HF's danger has nothing to do with acid strength. The fluoride ion penetrates skin, enters the bloodstream, and leaches calcium from blood — calcium that your heart needs to maintain a normal rhythm. Even small skin exposures warrant an emergency room visit. HF is extremely toxic, but it is not a strong acid.
Recognizing Weak Acids: Carboxylic Acids
Most acids you encounter that are not on the strong acid list are weak acids. One important category to recognize is carboxylic acids — organic acids containing the COOH (or CO2H) group. Acetic acid (CH3COOH) is the classic example.
Spotting Carboxylic Acids
Carboxylic acids are tricky because they do not begin with H in the formula — the acidic hydrogen is at the end, as part of the COOH group. Not every organic compound ending in H is an acid. Look specifically for the –COOH or –CO2H ending. If it has that group, it is a weak acid.
Weak Bases: Ammonia and Organic Amines
Ammonia (NH3) is the most recognizable weak base. Anything structurally resembling ammonia is likely a weak base as well. Organic amines are ammonia derivatives where one or more hydrogens have been replaced by carbon chains — for example, methylamine (CH3NH2).
Beyond ammonia and its derivatives, recognizing weak bases is tricky because there is no simple list to memorize. The good news is that most general chemistry courses will not expect you to identify weak bases beyond ammonia and amine-like structures.
Key Takeaways
- Three definitions exist for acids and bases — Arrhenius (aqueous only, H3O+/OH−), Brønsted-Lowry (proton donor/acceptor), and Lewis (electron pair acceptor/donor) — each progressively more comprehensive.
- Brønsted-Lowry is the most useful working definition: acids donate H+ and bases accept H+. Follow the proton to identify the acid and base in any reaction.
- Lewis has the ultimate definition — every acid recognized by any definition is guaranteed to be a Lewis acid. Uniquely Lewis acids include boron compounds (empty orbitals) and metal ions.
- A Lewis base must have a lone pair of electrons to donate. If that lone pair bonds to H+, the species is also a Brønsted-Lowry base.
- To find a conjugate base, remove H+ from the acid. To find a conjugate acid, add H+ to the base.
- Amphiprotic species (like HCO3−) can act as either acid or base, giving them both a conjugate acid and a conjugate base.
- Seven strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3) dissociate ~100% in water. Any acid not on this list is weak.
- Strong bases are Group 1 metal hydroxides and most Group 2 metal hydroxides. Group 2 hydroxides are strong but have limited solubility.
- HF is a weak acid despite being extremely dangerous — its toxicity comes from the fluoride ion, not acid strength.
- Weak acids and bases dissociate partially (often less than 5%). Key weak acids include carboxylic acids (COOH group); key weak bases include ammonia and organic amines.