AP Chemistry 8.1: Acid-Base Foundations, pH, and Unit 8 Review

Acid-base chemistry combines several ideas from earlier AP Chemistry units. You need equilibrium because weak acids and bases only partially ionize, stoichiometry because acids and bases react in specific mole ratios, and molecular structure because the arrangement of atoms in a molecule can affect acid strength.

• The most important goal of this unit is not simply memorizing which substances are acids and bases. You need to understand why acids and bases behave differently, how their strength is related to equilibrium, how to calculate pH, how buffers resist pH changes, and how titration curves reveal information about an unknown solution.
• A good way to see the entire unit is:
• acid/base identity → H₃O⁺ or OH⁻ concentration → pH → equilibrium → Ka/Kb → reactions → buffers → titrations

An acid is a substance capable of donating a proton, H⁺, while a base is a substance capable of accepting a proton. This is the Brønsted-Lowry definition of acids and bases.

• For example:
• HCl + H₂O → H₃O⁺ + Cl⁻
• HCl gives a proton to H₂O, so:
• HCl = acid
• H₂O = base
• Water accepts the proton and becomes H₃O⁺.
• In aqueous chemistry, free H⁺ is commonly represented as H₃O⁺, called the hydronium ion, because a proton interacts strongly with surrounding water molecules.

When an acid loses H⁺, the species remaining is its conjugate base.

When a base gains H⁺, the resulting species is its conjugate acid.

• Consider:
• NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
• NH₃ accepts H⁺, so NH₃ is the base.
• NH₄⁺ is its conjugate acid.
• H₂O donates H⁺, so H₂O is the acid.
• OH⁻ is its conjugate base.
• The pairs are therefore:
• NH₃ / NH₄⁺
• and:
• H₂O / OH⁻

Members of a conjugate acid-base pair differ by exactly one H⁺.

An easy way to identify them is:

• acid − H⁺ = conjugate base
• base + H⁺ = conjugate acid
• For example:
• H₂CO₃ → HCO₃⁻ → CO₃²⁻
• Removing one H⁺ from H₂CO₃ produces HCO₃⁻.
• Removing another H⁺ produces CO₃²⁻.

Some substances can behave as either acids or bases depending on what they react with.

• Water is the most important example.

With HCl:

• HCl + H₂O → H₃O⁺ + Cl⁻
• water accepts H⁺, so it acts as a base.

With NH₃:

• NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
• water donates H⁺, so it acts as an acid.

A substance capable of donating or accepting a proton is called amphiprotic.

Water can even react with itself:

• 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
• One water molecule donates H⁺ while another accepts it.

The equilibrium constant for this process is the ion-product constant of water:

• Kw = [H₃O⁺][OH⁻]
• At 25°C:
• Kw = 1.0 × 10⁻¹⁴

In pure water at 25°C:

• [H₃O⁺] = [OH⁻]
• Therefore:
• [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
• This is why pure water has a pH of 7 at 25°C.

An important detail is that Kw depends on temperature, so neutral pH is not universally exactly 7 at every temperature. Neutrality really means:

• [H₃O⁺] = [OH⁻]

At 25°C:

An acidic solution has:

• [H₃O⁺] > [OH⁻]

A neutral solution has:

• [H₃O⁺] = [OH⁻]

A basic solution has:

• [OH⁻] > [H₃O⁺]

You can use Kw to find one concentration from the other:

• [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴
• Suppose:
• [H₃O⁺] = 1.0 × 10⁻⁴ M
• Then:
• [OH⁻] = (1.0 × 10⁻¹⁴)/(1.0 × 10⁻⁴)
• [OH⁻] = 1.0 × 10⁻¹⁰ M
• Since hydronium concentration is much larger, the solution is acidic.

Instead of constantly writing extremely small H₃O⁺ concentrations, chemists use pH:

• pH = −log[H₃O⁺]
• Similarly:
• pOH = −log[OH⁻]
• At 25°C:
• pH + pOH = 14.00
• This relationship comes from Kw.

Suppose:

• [H₃O⁺] = 1.0 × 10⁻³ M
• Then:
• pH = −log(1.0 × 10⁻³)
• pH = 3.00
• The solution is acidic.

Suppose instead:

• [OH⁻] = 1.0 × 10⁻⁴ M
• Then:
• pOH = 4.00
• and:
• pH = 14.00 − 4.00
• pH = 10.00
• The solution is basic.

If pH is known:

• [H₃O⁺] = 10⁻ᵖᴴ

If pOH is known:

• [OH⁻] = 10⁻ᵖᴼᴴ
• For example:
• pH = 5.00
• [H₃O⁺] = 10⁻⁵
• [H₃O⁺] = 1.0 × 10⁻⁵ M

The pH scale is logarithmic, not linear.

• A solution with pH 3 has:
• 10 times greater [H₃O⁺] than pH 4
• and:
• 100 times greater [H₃O⁺] than pH 5
• Therefore, a change of one pH unit corresponds to a factor-of-10 change in hydronium concentration.

Do not assume an acid must have H written first in its formula. Acid-base behavior depends on proton transfer.

Do not assume a negatively charged species must be a base or a positively charged species must be an acid without examining the reaction.

Remember that conjugate acid-base pairs differ by exactly one H⁺.

At 25°C:

pH + pOH = 14

but this comes from the temperature-dependent value of Kw.

Acid = H⁺ donor

• Base = H⁺ acceptor
• Kw = [H₃O⁺][OH⁻]
• pH = −log[H₃O⁺]
• pOH = −log[OH⁻]

Acid/base definition: Acid donates H⁺; base accepts H⁺.

Conjugate pairs: They differ by exactly one H⁺.

Water: Can act as either an acid or base.

Neutrality: Means [H₃O⁺]=[OH⁻], not universally “pH must always equal 7.”

Strong vs. concentrated: They are different concepts.

Weak acids: Do not assume complete ionization.

Weak bases: Solve for OH⁻ first when using Kb.

Ka: Larger Ka means stronger acid.

pKa: Smaller pKa means stronger acid.

Conjugates: Strong acid has a very weak conjugate base.

Salt solutions: Ions from weak acids or weak bases can react with water.

Percent ionization: Dilution generally increases percent ionization of a weak acid even though [H₃O⁺] may decrease.

Structure: Compare conjugate-base stability rather than blindly using one periodic trend.

Acid-base mixtures: Stoichiometry comes before equilibrium.

Buffers: Must contain meaningful amounts of a weak species and its conjugate partner.

Henderson-Hasselbalch: Use base/acid, not acid/base.

Buffer capacity: Same pH does not mean same buffer capacity.

Titrations: Equivalence means stoichiometric equivalence, not equal volume.

Half-equivalence: For weak acid titrated with strong base, pH=pKa.

Weak-acid equivalence: Usually pH>7 with strong-base titrant.

Weak-base equivalence: Usually pH<7 with strong-acid titrant.

Indicators: Endpoint and equivalence point are not exactly the same concept.

Polyprotic acids: Can produce multiple equivalence points.

Before taking a Unit 8 test, you should be able to identify Brønsted-Lowry acids and bases, find conjugate acid-base pairs, and explain why water can behave as either an acid or a base.

You should understand water autoionization, use Kw, calculate pH and pOH, convert between pH and ion concentration, and understand that the pH scale is logarithmic.

You should be able to distinguish strong vs. weak from concentrated vs. dilute, calculate the pH of strong acid/base solutions, and account for stoichiometric numbers of H₃O⁺ or OH⁻ when appropriate.

For weak acids and bases, you should be able to write Ka and Kb expressions, construct ICE tables, calculate equilibrium concentrations and pH, calculate percent ionization, use small-x approximations appropriately, and understand the relationship:

• KaKb = Kw

You should be able to compare acid strengths using Ka and pKa, understand the acid-base behavior of salts, and recognize how conjugate acid/base strength is related.

You should be able to explain acid strength using molecular structure, especially bond strength, electronegativity, atomic size, resonance, and stabilization of the conjugate base.

For acid-base reactions, you should know when to perform stoichiometric neutralization before equilibrium calculations, calculate leftover strong acid or base, and identify the important species remaining after reaction.

For buffers, you should understand how a weak acid/conjugate-base pair resists changes in pH, use the Henderson-Hasselbalch equation, calculate the effect of added acid/base, understand buffer range, and distinguish buffer pH from buffer capacity.

Finally, you should be able to read and calculate from acid-base titration curves, including strong acid-strong base, weak acid-strong base, and weak base-strong acid titrations. You should know what chemistry controls pH in the initial region, buffer region, half-equivalence point, equivalence point, and after-equivalence region.

The biggest Unit 8 pattern to remember is:

• identify the species → determine whether reaction or equilibrium happens first → calculate amounts/concentrations → choose Ka, Kb, buffer, or strong-acid/base math → calculate pH
• Once you know what species are actually present at that moment, most Unit 8 problems become much easier.