AP Chemistry 8.5 Acid-Base Titrations Study Notes - New Syllabus Effective fall 2024

AP Chemistry 8.5 Acid-Base Titrations Study Notes- New syllabus

AP Chemistry 8.5 Acid-Base Titrations Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.

LEARNING OBJECTIVE

Explain results from the titration of a mono- or polyprotic acid or base solution, in relation to the properties of the solution and its components.

Key Concepts:

Titration Curves & Equivalence Points
Titration Curves & Half-Equivalence Points
Titration Curves of Polyprotic Acids

AP Chemistry-Concise Summary Notes- All Topics

8.5.A.1 Titration Curve:

1. Acids and Bases:

i. Strong Acids and Bases:

These fully dissociate in water, i.e., 100% of the acid or base molecules disintegrate into ions.

Strong Acids (examples):

* HCl → H⁺ + Cl⁻
* HNO₃ → H⁺ + NO₃⁻
* H₂SO₄ (first proton only) → H⁺ + HSO₄⁻

In water:

* All acid molecules give up protons (H⁺).
* No original acid molecules are left.
* The solution is highly conductive (lots of free ions).
* pH is extremely low (for acids) or extremely high (for bases).

Strong Bases (examples):

* NaOH → Na⁺ + OH⁻
* KOH → K⁺ + OH⁻
* Ba(OH)₂ → Ba²⁺ + 2OH⁻

In water:

* Fully release OH⁻ ions.
* High hydroxide ion concentration.

ii. Weak Acids and Bases:

These only partially dissociate in water. Most of the molecules stay intact.

Weak Acids (examples):

* CH₃COOH (acetic acid) ⇌ H⁺ + CH₃COO⁻
* HF (hydrofluoric acid) ⇌ H⁺ + F⁻

In water:

* Only a small percentage releases H⁺ ions.
* Forms an equilibrium (⇌).
* pH is higher (less acidic) than strong acids of the same concentration.

Weak Bases (examples):

* NH₃ (ammonia) + H₂O ⇌ NH₄⁺ + OH⁻
* CH₃NH₂ + H₂O ⇌ CH₃NH₃⁺ + OH⁻

In water:

* Few molecules only accept protons to become OH⁻.
* Also establishes equilibrium.

2. Titration Process:

i. Titration: What It Is:

Titration is a laboratory method of finding the concentration of an unknown solution (the analyte) by reacting it with a solution of known concentration (the titrant).

Key Terms:

Term	Definition
Titrant	The known solution added from a burette (usually an acid or base).
Analyte	The unknown solution being analyzed.
Equivalence Point	The point where moles of titrant = moles of analyte (neutralization).
End Point	The point where the indicator changes color, signaling the equivalence point (ideally very close).
Indicator	A dye (e.g., phenolphthalein) that changes color at or near the equivalence point.

ii. How Titration Works – Step-by-Step:

1. Prepare your solutions

* Titrant (known concentration) into the burette.
* Analyte (concentration is unknown) goes into a flask, usually a couple of drops of indicator.

2. Slowly add titrant:

* Apply the burette to add titrant to the analyte slowly.
* Stir the flask repeatedly so that solutions become mixed.

3. Look out for the end point

* While you slowly add the titrant, the indicator will change color.
* Stop adding when a permanent color change is achieved — this is the end point.

4. Determine the unknown concentration

* Apply the volume of titrant added and a balanced chemical equation to determine the moles of analyte, then compute its concentration.

iii. Titration Equation:

M₁V₁ = M₂V₂ (for monoprotic acid/base reactions)

Where:

* M₁, V₁ = molarity and volume of the acid
* M₂, V₂ = molarity and volume of the base

For polyprotic acids or more complicated reactions, apply stoichiometry using a balanced equation.

3. Titration Curve Interpretation:

i. A titration curve is a plot of pH (y-axis) against volume of titrant added (x-axis).

The form of the curve will be determined by the strength of the acid and base used.

ii. Key Parts of the Curve:

1. Initial pH (Starting Point):

* This is the pH prior to any addition of titrant.
* Varies with the strength of the analyte:

* Strong acid (e.g., HCl): extremely low pH (~1).
* Weak acid (e.g., CH₃COOH): higher pH (~3–4).
* Strong base: extremely high pH (~13–14).
* Weak base: lower (~9–10).

2. Gradual Rise or Fall:

* When you add titrant, the pH begins to change slowly.
* This area indicates the buffering capacity, particularly in weak acid/base titrations.

3. Steep Rise (or Fall):

* Abrupt, sharp pH change over a small volume.
* This is where the base and acid are quickly neutralizing one another.
* Happens close to the equivalence point.

4. Equivalence Point:

* Moles of titrant = moles of analyte.
* The solution is neutralized.
* The pH at equivalence varies depending on the type of titration:

Titration Type	pH at Equivalence
Strong acid + strong base	~7 (neutral)
Weak acid + strong base	>7 (slightly basic)
Strong acid + weak base	<7 (slightly acidic)

5. Beyond the Equivalence Point:

* Introducing more titrant overwhelms the analyte.
* pH reaches a plateau and indicates the type of the excess titrant:

* If titrant is a base → pH jumps high and becomes level.
* If titrant is an acid → pH drops low and becomes level.

8.5.A.2 Equivalence Point and Analyte Concentration:

1. Acid-Base Neutralization and Mole Concept:

i. Acid-Base Neutralization Reaction:

Neutralization is the reaction between an acid and a base to produce water and a salt:

ii. General Reaction (Monoprotic Case): $\text{Acid (H⁺ donor)} + \text{Base (OH⁻ donor)} → \text{Salt} + \text{Water}$

Example: $\text{HCl} + \text{NaOH} → \text{NaCl} + \text{H₂O}$

ii. Mole Concept in 1:1 Neutralization:

For monoprotic acids (donate one H⁺) and monobasic bases (donate one OH⁻), neutralization is a 1:1 mole ratio.

That is:

* 1 mole of H⁺ from the acid combines with 1 mole of OH⁻ from the base.
* Neutralization is complete when: $\text{moles of acid (H⁺)} = \text{moles of base (OH⁻)}$

2. Titration Calculations:

Apply the formula:

M1 V1 = M2 V2

Where:

* M1 = molarity of acid
* V1 = volume of acid
* M2 = molarity of base
* V2 = volume of base

> Applies to 1:1 mole ratio titrations (e.g., monoprotic acids and monobasic bases)

8.5.A.3 Half-Equivalence Point and pKa Determination:

1. Conjugate Acid-Base Pair:

During an acid-base reaction, when a weak acid (such as HA) loses a proton (H⁺), it becomes its conjugate base (A⁻). This is because acids are proton donors.

Example:
HA ⇌ H⁺ + A⁻

Here:

* HA is the weak acid.
* A⁻ is the conjugate base, which is formed after HA loses a proton.

In reverse, A⁻ may take up a proton to form HA once again, thus a base in the conjugate pair.

This acid-base conjugate pair (HA and A⁻) are connected as they only vary by one proton (H⁺). This is fundamental when interpreting buffer systems and titration curves, particularly at the half-equivalence point, where [HA] = [A⁻].

2. Henderson-Hasselbalch Equation: $pH = pKa + \log (\frac{[A^{-}]}{[H A]})$

* pH: Solution pH
* pKa: Acid dissociation constant
* [A⁻]: Concentration of the conjugate base
* [HA]: Concentration of weak acid

Key Point:

At the half-equivalence point (where

[A^-] = [HA]

), pH = pKa, allowing you to determine the pKa of a weak acid.

3. Half-Equivalence Point:

At the half-equivalence point, half of the weak acid (HA) has been transformed into its conjugate base (A⁻), so:

$[H A] = [A^{-}]$

This reduces the Henderson-Hasselbalch equation to: $pH = pKa$

So, at this point, the pH is equal to the pKa of the acid, and the pKa can be directly calculated from the measured pH.

4. Why It Matters:

By measuring the pH at the half-equivalence point in a titration, you can determine the pKa (and therefore Ka) of a weak acid. This is because, at this point, pH = pKa, providing a direct method for experimentally determining the acid’s dissociation constant.

8.5.A.5 Titration of Polyprotic Acids and pKa Determination:

1. Polyprotic Acids and Their Protons:

Polyprotic acids are able to donate multiple protons and experience sequential ionization steps, each with its own dissociation constant Ka.

i. Ionization Steps:

a. First ionization is stronger, with a larger Ka1

Example:

$\text{H}_2\text{SO}_4 \to \text{H}^+ + \text{HSO}_4^-$

(complete dissociation)

b. Second ionization is weaker, with a smaller Ka2.

Example:

$\text{HSO}_4^- \to \text{H}^+ + \text{SO}_4^{2-}$

(partial dissociation)

ii. Ka and pKa:

${pKa}_{1} < {pKa}_{2} < {pKa}_{3}$

2. Species at Different Points:

i. Before Equivalence (Initial Point):

* Weak acid titrated with strong base:

* Mostly acid (HA): Weak acid in excess.
* Conjugate base (A⁻): Being formed in minute quantities.
* H⁺ ions: In lower concentration as the acid is dissociating.
* Strong acid titrated with strong base:

* Predominantly H⁺ (from acid).
* OH⁻ ions: Added as the base is being titrated.

ii. At Half-Equivalence Point (Buffer Region)

* Weak acid titrated with strong base:

* [HA] = [A⁻]: Equivalents concentrations of the acid and its conjugate base.
* pH = pKa at this point.

* Strong acid titrated with strong base:

* Near neutral: Most of the acid has reacted and approximately equal volumes of acid and base remain.

iii. At Equivalence Point:

* Weak acid titrated with strong base:

* Conjugate base (A⁻) is in excess (because it arises from acid neutralization).
* Water and salt: The resultant salt (e.g., NaA) might hydrolyze, so the solution becomes somewhat basic (pH > 7).

* Strong acid titrated with strong base:

* Water and salt: The solution is neutral (pH = 7), since acid and base completely neutralize.

iv. Beyond Equivalence (Excess Titrant):

* Weak acid titrated with strong base:

* Excess base (OH⁻) is in excess, raising the pH.
* Strong acid titrated with strong base:

* Excess OH⁻: Adding more base drives the pH up.

8.5.A.4 Equivalence Point and pH in Titrations:

1. Strong vs. Weak Acid/Base Titrations:

Titration Type	pH at Equivalence	Why
Strong Acid + Strong Base	pH = 7	Neutral solution due to complete dissociation.
Weak Acid + Strong Base	pH > 7 (basic)	Conjugate base of weak acid hydrolyzes to form OH⁻.
Strong Acid + Weak Base	pH < 7 (acidic)	Conjugate acid of weak base hydrolyzes to form H⁺.

2. Conjugate Base/Acid Behavior:

i. Conjugate Base/Acid Behavior – Hydrolysis:

When an acid or base dissociates, it produces a conjugate species, which can react with water (hydrolyze) to either release H⁺ ions (creating an acidic solution) or accept H⁺ ions (creating a basic solution). This comes into play when thinking about the pH at the equivalence point in a titration.

ii. Conjugate Base (from Weak Acid):

* When a weak acid (e.g., acetic acid, CH₃COOH) is titrated with a strong base, the formed conjugate base (A⁻) will react with water:

$\text{A}⁻ + \text{H}_2\text{O} \rightleftharpoons \text{HA} + \text{OH}⁻$

* Conjugate base (A⁻) accepts H⁺ ions from water, forming OH⁻ ions, thereby making the solution basic.
* Result: pH at equivalence is greater than 7.

Example:
In the titration of acetic acid (CH₃COOH) with sodium hydroxide (NaOH), the acetate ion (CH₃COO⁻) is the conjugate base and will increase the pH.

iii. Conjugate Acid (from Weak Base)

* When a weak base (e.g., ammonia, NH₃) is titrated with a strong acid, the conjugate acid (e.g., NH₄⁺) produced can react with water:

$\text{NH}_4^+ + \text{H}_2\text{O} \rightleftharpoons \text{NH}_3 + \text{H}_3\text{O}^+$

* The conjugate acid (NH₄⁺) donates H⁺ ions to water, hence the solution acidic.
* Result: pH at equivalence is less than 7.

Example:
In ammonia (NH₃) titration with hydrochloric acid (HCl), the ammonium ion (NH₄⁺) is the conjugate acid and will decrease the pH.

AP Chemistry 8.5 Acid-Base Titrations Study Notes

AP Chemistry 8.5 Acid-Base Titrations Study Notes - New Syllabus Effective fall 2024

AP Chemistry 8.5 Acid-Base Titrations Study Notes- New syllabus

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