8.1 Introduction to Acids and Bases

Water naturally splits into tiny amounts of hydronium, $H_3O^+$, and hydroxide, $OH^-$, and this autoionization is described by $K_w=[H_3O^+][OH^-]=1.0\times10^{-14}$ at 25 degrees C. You report those ion concentrations as $\text{pH}=-\log[H_3O^+]$ and $\text{pOH}=-\log[OH^-]$, and in pure water at 25 degrees C, $\text{pH}=\text{pOH}=7.0$. For AP Chemistry, remember that neutral means $[H_3O^+]=[OH^-]$.

Introduction to Acids and Bases Summary

AP Chemistry Topic 8.1 starts acids and bases with water autoionization and the pH scale. Hydronium concentration is reported as pH, hydroxide concentration is reported as pOH, and at 25 degrees C, Kw = [H3O+][OH-] = 1.0 x 10^-14.

At 25 degrees C, pH + pOH = 14.0 and pure water is neutral because pH = pOH = 7.0. At other temperatures, Kw changes, so neutral water can have a pH that is not exactly 7.0. The key is that neutral means [H3O+] = [OH-], not always pH 7.

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Why This Matters for the AP Chemistry Exam

This topic gives you the foundation for everything in the acids and bases unit, which carries a large share of the exam. Once you can move between [H3O+], [OH-], pH, and pOH using Kw, you are ready to handle strong acids and bases, weak acid and base equilibria, buffers, and titration calculations later in the unit.

You will use these relationships to solve calculation-based multiple-choice questions and to support claims with numbers and reasoning on free-response questions. The exam expects you to pick the right relationship for a problem and carry units and significant figures correctly.

Key Takeaways

• pH = -log[H3O+] and pOH = -log[OH-]; you can reverse them with [H3O+] = 10^(-pH) and [OH-] = 10^(-pOH).
• Water autoionizes: Kw = [H3O+][OH-] = 1.0 x 10^-14 at 25 degrees C.
• At 25 degrees C, pH + pOH = pKw = 14.0, so if you know one, you can find the other.
• A neutral solution has pH = pOH; in pure water at 25 degrees C that means pH = pOH = 7.0.
• Kw depends on temperature, so the pH of neutral water shifts away from 7.0 at temperatures other than 25 degrees C.
• "Hydrogen ion" (H+) and "hydronium ion" (H3O+) refer to the same thing in water; H3O+ is preferred, but H+ is accepted on the exam.

Acids and Bases: Two Common Definitions

There are two definitions you should know before working with pH.

Arrhenius Definition

The Arrhenius definition calls an acid any compound that increases the concentration of hydrogen ions (H+) in solution, and a base any compound that increases the concentration of hydroxide ions (OH-). In short, an Arrhenius acid yields H+ in water, and an Arrhenius base yields OH-.

For example, HCl (hydrochloric acid) acts as an Arrhenius acid because, in water:

HCl -> H+ + Cl-

HCl yields an H+ ion in water, which makes it an Arrhenius acid.

Bronsted-Lowry Definition

The Bronsted-Lowry definition treats acids and bases in terms of proton transfer. An acid donates an H+ ion, and a base accepts it:

HA + B- -> HB + A-

Here HA is the acid, which donates an H+ ion to B- (the base) to form HB and A-.

Get used to seeing HA and B- as sample acids and bases; it is common notation. A and B can stand in for many different ions.

Many of these reactions are reversible, so you will often see them written as an equilibrium:

HA + B- ⇄ HB + A-

The Hydronium Ion

When a Bronsted acid dissolves in water, it needs something to donate its H+ to. It donates to water, forming H3O+ (hydronium). So you can show an acid dissolving in an Arrhenius way as HA ⇄ H+ + A- and in a Bronsted way as HA + H2O ⇄ H3O+ + A-. These mean the same thing, and they show the difference between the two definitions.

You will see both "hydrogen ion" and "hydronium ion" used, along with the symbols H+(aq) and H3O+(aq). They are often used interchangeably because they represent the same idea: a proton in water. H3O+(aq) is more accurate since bare protons do not exist in water, but both H+(aq) and H3O+(aq) are accepted on the AP Chemistry exam.

Conjugate Acids and Bases

How do you classify HB and A- in that reaction? If the reaction ran in reverse, A- would act as a base and HB would act as an acid. They get special names: HB is the conjugate acid of B-, and A- is the conjugate base of HA.

When HA is a weak acid and B- is a weak base, their conjugates (A- and HB) have meaningful acidity or basicity. The conjugate of a strong acid or strong base has little to no acidity or basicity. The general pattern: the weaker the acid or base, the stronger its conjugate.

Autoionization of Water and Kw

Water autoionizes, meaning a proton transfers from one water molecule to another, producing a hydronium ion (H3O+) and a hydroxide ion (OH-). This equilibrium has its own constant, Kw:

Kw = [H3O+][OH-] = 1.0 x 10^-14 at 25 degrees C

In pure water, [H3O+] = [OH-] = 1.0 x 10^-7 M at 25 degrees C, so pH = pOH. A solution where pH = pOH is called neutral.

Because Kw = 1.0 x 10^-14 at 25 degrees C, taking the negative log of both sides gives:

pKw = 14.0 = pH + pOH at 25 degrees C

So in any aqueous solution at 25 degrees C, pH + pOH = 14.0. If you know one value, subtract from 14.0 to get the other.

Temperature Changes Kw

Kw is temperature dependent. Autoionization is the process being measured, and raising or lowering the temperature shifts how much water autoionizes, which changes Kw. That means the pH of pure, neutral water deviates from 7.0 at temperatures other than 25 degrees C. Even when pH is not 7.0, pure water is still neutral because pH still equals pOH.

Calculating pH, pOH, and Ion Concentrations

The concentrations of hydronium and hydroxide ions are reported as pH and pOH:

• pH = -log[H3O+]
• pOH = -log[OH-]

To go the other direction:

• [H3O+] = 10^(-pH)
• [OH-] = 10^(-pOH)

Worked Example

Suppose [H3O+] = 1.0 x 10^-5 M.

pH = -log(1.0 x 10^-5) = 5.0

Use Kw to find [OH-]:

[OH-] = Kw / [H3O+] = (1.0 x 10^-14) / (1.0 x 10^-5) = 1.0 x 10^-9 M

Then:

pOH = -log(1.0 x 10^-9) = 9.0

Check: pH + pOH = 5.0 + 9.0 = 14.0, which matches pKw at 25 degrees C.

How to Use This on the AP Chemistry Exam

Problem Solving

• Identify which quantity you are given (concentration, pH, or pOH) and choose the matching relationship.
• Use [H3O+] = 10^(-pH) and [OH-] = 10^(-pOH) to move from a log value back to a concentration.
• Use Kw = [H3O+][OH-] to connect the two ion concentrations when you only have one.
• Use pH + pOH = 14.0 only at 25 degrees C as a fast shortcut.

Significant Figures

• In a logarithm, only the digits after the decimal point count as significant figures. So [H3O+] = 1.0 x 10^-5 M (two sig figs) gives pH = 5.00, where the "00" are the significant digits.
• Carry units on concentrations (M) and keep them out of pH and pOH, which are unitless.

Common Trap

• Do not assume neutral always means pH 7. Neutral means pH = pOH. At temperatures other than 25 degrees C, neutral water has a pH other than 7.0.

Common Misconceptions

• "Neutral always equals pH 7." Neutral means pH = pOH. That equals 7.0 only at 25 degrees C, because Kw changes with temperature.
• "pH 7 is required for neutral." Pure water stays neutral at other temperatures even though its pH is not exactly 7.0.
• "H+ and H3O+ are different species you must keep separate." In water they represent the same thing; H3O+ is preferred, but H+ is accepted on the exam.
• "Kw is a fixed number." Kw is 1.0 x 10^-14 only at 25 degrees C; it shifts with temperature.
• "A strong acid has a strong conjugate base." The conjugate of a strong acid or base has little to no acid or base strength; weaker acids and bases have stronger conjugates.
• "pH always falls between 0 and 14." That range is a common scale, not a hard limit. Concentrated solutions can give pH values outside 0 to 14.

Related AP Chemistry Guides

• Unit 8 Overview: Acids and Bases
• 8.2 pH and pOH of Strong Acids and Bases
• 8.4 Acid-Base Reactions and Buffers
• 8.3 Weak Acid and Base Equilibria
• 8.5 Acid-Base Titrations
• 8.7 pH and pKa