🎯⭐ INTERACTIVE LESSON

Acid-Base Theories and pH Scale

Learn step-by-step with interactive practice!

← Back to Standard Lesson

Acid-Base Theories and pH Scale - Complete Interactive Lesson

Part 1: Arrhenius & Brønsted-Lowry

🧪 Arrhenius Acids and Bases

Part 1 of 7 — The First Modern Definition

Three Acid-Base Theories — Where We're Headed

Theory	Acid Is...	Base Is...	This Part
Arrhenius	Produces $H^+$ in water	Produces $OH^-$ in water	✅ Part 1
Brønsted-Lowry	Proton donor	Proton acceptor	Part 2
Lewis	Electron pair acceptor	Electron pair donor	Part 3

🔑 Why this matters: The Arrhenius model is the foundation — every acid-base theory that follows builds on these ideas.

What You'll Master in Part 1

Defining Arrhenius acids and bases by what they produce in water
Identifying limitations of the Arrhenius model
Recognizing strong acids and the hydronium ion concept

📖 The Arrhenius Definition

In 1884, Svante Arrhenius proposed a simple classification:

Type	Definition	Example
Arrhenius Acid	Produces $H^+$ ions in aqueous solution	$HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$

🧪 Common Arrhenius Acids

Strong Acids (Complete Dissociation)

Formula	Name	Dissociation
$HCl$	Hydrochloric acid	$HCl \rightarrow H^+ + Cl^-$

⚛️ The Hydronium Ion

In reality, free $H^+$ ions (bare protons) don't exist in water. Instead, they bond to water molecules:

$H^+(aq) + H_2O(l) \rightarrow H_3O^+(aq)$

Arrhenius Concept Check 🎯

📌 Arrhenius Neutralization

When an Arrhenius acid reacts with an Arrhenius base, they undergo neutralization:

$\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

Examples

$HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$

Arrhenius Classification 🔍

Exit Quiz — Arrhenius Acids & Bases ✅

Part 2: Conjugate Acid-Base Pairs

🔄 Brønsted-Lowry Acids and Bases

Part 2 of 7 — Proton Donors and Acceptors

Arrhenius vs. Brønsted-Lowry

Feature	Arrhenius	Brønsted-Lowry
Acid definition	Produces $H^+$ in water	Donates a proton ( $H^+$ )
Base definition	Produces in water

Part 3: The pH Scale

🔬 Lewis Acids and Bases

Part 3 of 7 — Electron Pair Donors and Acceptors

How the Three Theories Compare

Theory	Key Question	Broadest?
Arrhenius	Does it produce $H^+$ or $OH^-$ ?	Narrowest
Brønsted-Lowry

Part 4: Strong Acids & Bases

📊 The pH Scale

Part 4 of 7 — Measuring Acidity and Basicity

The pH Scale at a Glance

pH	$[H^+]$ (M)	Character	Example
0	$10^0$	Strongly acidic

Part 5: Calculating pH & pOH

💪 Strong Acids and Bases — pH Calculations

Part 5 of 7 — Complete Dissociation Means Easy Math

Strong = Complete Dissociation

Type	Example	Key Calculation
Strong monoprotic acid	0.025 M HCl	$[H^+] = 0.025$ M, pH = 1.60
Strong diprotic acid	0.010 M $H_2SO_4$

Part 6: Problem-Solving Workshop

🛠️ Problem-Solving Workshop

Part 6 of 7 — Acid-Base Theories and pH

Problem Types You'll Practice

Problem Type	Skills Combined
Theory identification	Arrhenius vs. Brønsted-Lowry vs. Lewis
Multi-step pH	Dilution → dissociation → $-\log$
Conceptual reasoning	Very dilute acid pH limits
Conjugate pair analysis	Identifying donors/acceptors

🔑 Why this matters: AP Chemistry free-response questions often combine acid-base theory with pH calculations — exactly the type of multi-step problems in this workshop.

What You'll Master in Part 6

Solving multi-step pH problems with dilution
Identifying acid-base behavior across all three theories
Reasoning about edge cases like very dilute strong acids

🧪 Problem 1: Identifying Acid-Base Behavior

Consider these reactions:

Part 7: Synthesis & AP Review

🎓 Synthesis & AP Review

Part 7 of 7 — Acid-Base Theories and pH

Everything Comes Together

Topic	Key Equation or Concept
Three theories	Arrhenius ⊂ Brønsted-Lowry ⊂ Lewis
pH/pOH	$pH + pOH = 14$
Strong acids	$[H^+]$ = concentration (complete dissociation)

Acids increase $[H^+]$ in water
Bases increase $[OH^-]$ in water
Neutralization produces water: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$

🔑 Key idea: Arrhenius acids add $H^+$ to solution; Arrhenius bases add $OH^-$ .

The Arrhenius model only works in aqueous solutions and cannot explain:

Why $NH_3$ acts as a base (it doesn't contain $OH^-$ )
Acid-base behavior in non-aqueous solvents
Reactions between gases that show acid-base character

⚠️ These limitations led to the development of the broader Brønsted-Lowry and Lewis definitions (Parts 2–3).

Formula	Name	Dissociation
$NaOH$	Sodium hydroxide	$NaOH \rightarrow Na^+ + OH^-$
$KOH$	Potassium hydroxide	$KOH \rightarrow K^+ + OH^-$
$Ca(OH)_2$	Calcium hydroxide	$Ca(OH)_2 \rightarrow Ca^{2+} + 2OH^-$
$Ba(OH)_2$	Barium hydroxide	$Ba(OH)_2 \rightarrow Ba^{2+} + 2OH^-$

💡 Memorize the 6 strong acids and 4 strong bases — everything else is weak!

The hydronium ion $H_3O^+$ is a more accurate representation. In AP Chemistry:

$H^+(aq)$ and $H_3O^+(aq)$ are used interchangeably
Both notations are acceptable on the AP exam
$H_3O^+$ is technically more correct
$H^+$ is a convenient shorthand

Autoionization of Water

Pure water undergoes self-ionization:

$2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$

The equilibrium constant for this process is:

$\boxed{K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \text{ at 25°C}}$

🔑 In pure water: $[H^+] = [OH^-] = 1.0 \times 10^{-7}$ M

H Cl (a q) + N a O H (a q) \to N a Cl (a q) + H_{2} O (l)

Net ionic equation:

$\boxed{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}$

🔑 This net ionic equation is the same for all strong acid–strong base neutralizations!

Double Replacement Pattern

$H_2SO_4(aq) + 2KOH(aq) \rightarrow K_2SO_4(aq) + 2H_2O(l)$

💡 Sulfuric acid is diprotic — it has 2 acidic protons, so it requires 2 moles of $KOH$ .

🔑 Why this matters: The Brønsted-Lowry model is what the AP exam uses most — conjugate acid-base pairs appear in nearly every acid-base question.

What You'll Master in Part 2

Identifying proton donors (acids) and proton acceptors (bases)
Writing conjugate acid-base pairs for any reaction
Recognizing amphoteric substances like water

📖 The Brønsted-Lowry Definition

Type	Definition
Brønsted-Lowry Acid	A proton ( $H^+$ ) donor
Brønsted-Lowry Base	A proton ( $H^+$ ) acceptor

Key Advantage

This definition works in any solvent — not just water!

🔑 Unlike Arrhenius, Brønsted-Lowry doesn't require water — any proton transfer counts.

Example: $HCl$ in Water

$HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$

$HCl$ donates a proton → acid
$H_2O$ accepts a proton → base

Example: $NH_3$ in Water

$NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$

$NH_3$ accepts a proton → base
$H_2O$ donates a proton → acid

💡 Water can act as either an acid or a base! This is called being amphoteric (or amphiprotic).

🧪 Conjugate Acid-Base Pairs

When an acid donates a proton, the product is its conjugate base. When a base accepts a proton, the product is its conjugate acid.

$\underbrace{HA}_{\text{acid}} + \underbrace{B}_{\text{base}} \rightleftharpoons \underbrace{A^-}_{\text{conjugate base}} + \underbrace{BH^+}_{\text{conjugate acid}}$

Examples

Acid	Conjugate Base	Relationship
$HCl$	$Cl^-$	Differs by one $H^+$

Critical Rule

🔑 A conjugate pair always differs by exactly one proton ( $H^+$ ).

Strength Relationship

⚠️ Strong acid → very weak conjugate base (and vice versa)

$HCl$ is strong → $Cl^-$ is a negligible base (does not accept protons)
$CH_3COOH$ is weak → is a moderate conjugate base

Brønsted-Lowry Concept Check 🎯

⚗️ Identifying Conjugate Pairs in Reactions

For any Brønsted-Lowry reaction, there are always two conjugate pairs:

$\underbrace{HF}_{\text{acid}_1} + \underbrace{H_2O}_{\text{base}_2} \rightleftharpoons \underbrace{F^-}_{\text{conj. base}_1} + \underbrace{H_3O^+}_{\text{conj. acid}_2}$

Pair 1: $HF / F^-$

Pair 2: $H_2O / H_3O^+$

Steps to Identify

Find the species that lost a proton → that's the acid; its product is the conjugate base
Find the species that gained a proton → that's the base; its product is the conjugate acid
Each acid is paired with its conjugate base (they differ by one $H^+$ )

Conjugate Pair Identification 🔍

For the reaction: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

Conjugate Pair Practice 🧮

Identify the conjugate partners:

1) What is the conjugate base of $H_2CO_3$ ? (Enter the chemical formula, e.g. Cl-)

2) What is the conjugate acid of $PO_4^{3-}$ ? (Enter the chemical formula, e.g. H2SO4)

3) What is the conjugate base of $H_2O$ ? (Enter the chemical formula, e.g. F-)

Exit Quiz — Brønsted-Lowry Theory ✅

The Lewis model captures reactions that have nothing to do with protons!

🔑 Why this matters: Lewis acid-base theory explains coordination chemistry, organic reactions, and metal complex formation — all tested on the AP exam.

What You'll Master in Part 3

Defining Lewis acids (electron pair acceptors) and bases (electron pair donors)
Identifying Lewis acids: metal cations, incomplete octets, $H^+$
Comparing all three acid-base theories on the AP exam

📖 The Lewis Definition

Type	Definition	Key Feature
Lewis Acid	Electron pair acceptor	Has an empty orbital or can make room for electrons
Lewis Base	Electron pair donor	Has a lone pair of electrons to share

Comparison of All Three Theories

Theory	Acid	Base
Arrhenius	Produces $H^+$ in water	Produces $OH^-$ in water
Brønsted-Lowry	Proton donor	Proton acceptor
Lewis	Electron pair acceptor	Electron pair donor

Key Insight

Every Arrhenius acid is a Brønsted-Lowry acid, and every Brønsted-Lowry acid involves a Lewis acid interaction. The Lewis definition is the most inclusive.

$\boxed{\text{Arrhenius} \subset \text{Br\o nsted-Lowry} \subset \text{Lewis}}$

🔑 If you can’t explain a reaction with Arrhenius or Brønsted-Lowry, try Lewis — it covers everything.

🧪 Common Lewis Acids

1. Metal Cations

Metal ions have empty orbitals and accept electron pairs from ligands:

$Cu^{2+} + 4NH_3 \rightarrow [Cu(NH_3)_4]^{2+}$

$Cu^{2+}$ : Lewis acid (accepts electron pairs)
$NH_3$ : Lewis base (donates lone pair)

2. Molecules with Incomplete Octets

$BF_3 + NH_3 \rightarrow F_3B\text{-}NH_3$

$BF_3$ : Lewis acid (boron has only 6 electrons, empty p orbital)
$NH_3$ : Lewis base (nitrogen has a lone pair)

💡 Molecules with incomplete octets (like $BF_3$ and $AlCl_3$ ) are classic Lewis acids.

3. Protons ( $H^+$ )

The proton itself is a Lewis acid — it accepts an electron pair:

$H^+ + OH^- \rightarrow H_2O$

This shows how the Lewis definition encompasses the Brønsted-Lowry definition.

✏️ Common Lewis Bases

🔑 Any species with a lone pair can be a Lewis base:

$NH_3$ , $H_2O$ , $OH^-$ , ,

Lewis Acid-Base Concept Check 🎯

🔗 Coordinate Covalent Bonds

When a Lewis base donates an electron pair to a Lewis acid, the resulting bond is called a coordinate covalent bond (or dative bond).

$F_3B + :NH_3 \rightarrow F_3B\text{←}NH_3$

The arrow ← shows that both electrons in the bond came from the nitrogen of $NH_3$ .

💡 A coordinate covalent bond (dative bond) is formed whenever a Lewis base donates a lone pair to a Lewis acid.

In Coordination Chemistry

Metal ions form coordination compounds with Lewis bases (called ligands):

$Fe^{3+} + 6CN^- \rightarrow [Fe(CN)_6]^{3-}$

Lewis Acid	Lewis Base (Ligand)	Product
$Fe^{3+}$	$CN^-$	$[F e (C N_{}$

Lewis Acid-Base Classification 🔍

Theory Comparison 🧮

For each species, identify which acid-base theory can explain its behavior as an acid or base:

1) $NaOH$ acting as a base — which is the simplest theory that explains this? (Enter: Arrhenius, Bronsted-Lowry, or Lewis)

2) $NH_3$ acting as a base (no $OH^-$ in its formula) — simplest theory? (Enter: Arrhenius, Bronsted-Lowry, or Lewis)

3) $BF_3$ acting as an acid (no $H^+$ to donate) — simplest theory? (Enter: Arrhenius, Bronsted-Lowry, or Lewis)

Exit Quiz — Lewis Acids & Bases ✅

Each pH unit = a 10-fold change in $[H^+]$ .

🔑 Why this matters: pH calculations are on virtually every AP Chemistry exam — mastering the logarithmic relationship between $[H^+]$ and pH is essential.

What You'll Master in Part 4

Converting between $[H^+]$ , pH, $[OH^-]$ , and pOH
Understanding the pH + pOH = 14 relationship
Interpreting what pH values mean for acidity and basicity

🔗 pH, pOH, and Their Relationship

pH Definition

$\boxed{pH = -\log[H^+]}$

pOH Definition

$\boxed{pOH = -\log[OH^-]}$

The Key Relationship

At 25°C:

$\boxed{pH + pOH = 14}$

This comes from $K_w$ :

$[H^+][OH^-] = 1.0 \times 10^{-14}$

Taking $-\log$ of both sides:

$-\log[H^+] + (-\log[OH^-]) = -\log(1.0 \times 10^{-14})$

$pH + pOH = 14$

Interpreting pH

pH Range	Solution Type	$[H^+]$ vs $[OH^-]$
$pH < 7$

🔢 pH Calculations

From $[H^+]$ to pH

Problem: $[H^+] = 3.2 \times 10^{-4}$ M. Find pH.

Solution:

$pH = -\log(3.2 \times 10^{-4}) = -(-3.49) = 3.49$

From pH to $[H^+]$

Problem: $pH = 5.60$ . Find $[H^+]$ .

Solution:

$[H^+] = 10^{-pH} = 10^{-5.60} = 2.5 \times 10^{-6} \text{ M}$

From $[OH^-]$ to pH

Problem: $[OH^-] = 4.0 \times 10^{-3}$ M. Find pH.

Solution:

Step 1: $pOH = -\log(4.0 \times 10^{-3}) = 2.40$

Step 2: $pH = 14 - pOH = 14 - 2.40 = 11.60$

The "p" Notation

🔑 The prefix "p" always means $-\log$ :

$\boxed{pX = -\log X}$

So $pK_a = -\log K_a$ , $pK_b = -\log K_b$ ,

pH Concept Check 🎯

📌 Significant Figures in pH

An important AP Chemistry rule:

🔑 The number of decimal places in the pH equals the number of significant figures in $[H^+]$ .

Examples

$[H^+]$	Sig Figs	pH	Decimal Places
$1.0 \times 10^{-4}$	2	4.00	2
$2.5 \times 10^{-6}$

⚠️ The digits before the decimal in pH only indicate the order of magnitude — they don't count as sig figs!

pH Calculation Drill 🧮

1) What is the pH of a solution with $[H^+] = 5.0 \times 10^{-9}$ M? (2 decimal places)

2) What is the $[H^+]$ in a solution with $pH = 4.30$ ? (Enter in scientific notation, e.g. 2.5e-3)

3) What is the pH of a solution with $[OH^-] = 2.0 \times 10^{-4}$ M? (2 decimal places)

pH Scale Understanding 🔍

Exit Quiz — pH Scale ✅

🔑 Why this matters: Strong acid/base pH problems are the foundation for all later calculations — titrations, buffers, and equilibrium all build from here.

What You'll Master in Part 5

Calculating pH of strong monoprotic and diprotic acids
Calculating pOH and pH of strong bases (Group 1 and Group 2)
Handling dilution before calculating pH

🧪 pH of Strong Acids

For a strong acid $HA$ at concentration $C$ :

$HA \rightarrow H^+ + A^-$

Since dissociation is 100% complete: $[H^+] = C$

$\boxed{pH = -\log C}$

Example 1

Problem: What is the pH of 0.025 M $HCl$ ?

Solution:

$[H^+] = 0.025 \text{ M}$ $pH = -\log(0.025) = 1.60$

Example 2

Problem: What is the pH of 0.0040 M $HNO_3$ ?

Solution:

$[H^+] = 0.0040 \text{ M}$ $pH = -\log(0.0040) = 2.40$

Diprotic Strong Acid ( $H_2SO_4$ )

For the first dissociation (strong): $H_2SO_4 \rightarrow H^+ + HSO_4^-$

For dilute solutions, each mole of $H_2SO_4$ produces approximately 2 moles of $H^+$ :

$[H^+] \approx 2C \text{ (for dilute solutions)}$

⚠️ The second dissociation of $H_2SO_4$ ( $HSO_4^-$ ) is weak (), so at higher concentrations the approximation may not hold exactly.

📌 pH of Strong Bases

For a strong base like $NaOH$ at concentration $C$ :

$NaOH \rightarrow Na^+ + OH^-$

$[OH^-] = C$ , then:

$\boxed{pOH = -\log C \quad\Rightarrow\quad pH = 14 - pOH}$

Example 1

Problem: What is the pH of 0.010 M $NaOH$ ?

Solution:

$[OH^-] = 0.010 \text{ M}$ $pOH = -\log(0.010) = 2.00$

Group 2 Hydroxides

For $Ba(OH)_2$ or $Ca(OH)_2$ :

$Ba(OH)_2 \rightarrow Ba^{2+} + 2OH^-$

$[OH^-] = 2C$

Example 2

Problem: What is the pH of 0.0050 M $Ba(OH)_2$ ?

Solution:

$[OH^-] = 2(0.0050) = 0.010 \text{ M}$ $pOH = -\log(0.010) = 2.00$

Strong Acid/Base pH Check 🎯

📌 Mixing and Dilution

Diluting a Strong Acid

🔑 Use $M_1V_1 = M_2V_2$ for dilution problems:

Example: 25.0 mL of 0.10 M $HCl$ is diluted to 100.0 mL. What is the new pH?

$M_2 = \frac{M_1V_1}{V_2} = \frac{(0.10)(25.0)}{100.0} = 0.025 \text{ M}$

$pH = -\log(0.025) = 1.60$

Mixing Strong Acid and Strong Base

Example: 50.0 mL of 0.10 M $HCl$ + 30.0 mL of 0.10 M $NaOH$

Moles $H^+$ = $0.050 \times 0.10 = 0.0050$ mol

Moles $OH^-$ = $0.030 \times 0.10 = 0.0030$ mol

Excess $H^+$ = $0.0050 - 0.0030 = 0.0020$ mol

Total volume = $80.0$ mL = $0.0800$ L

$[H^+] = \frac{0.0020}{0.0800} = 0.025 \text{ M}$

$pH = -\log(0.025) = 1.60$

Strong Acid/Base Calculation Drill 🧮

1) What is the pH of 0.0020 M $HClO_4$ ? (2 decimal places)

2) What is the pH of 0.050 M $Ba(OH)_2$ ? (2 decimal places)

3) 40.0 mL of 0.15 M $HNO_3$ is mixed with 20.0 mL of 0.15 M $NaOH$ . What is the pH? (2 decimal places)

Strong Acid/Base Reasoning 🔍

Exit Quiz — Strong Acid/Base pH ✅

HF(aq) + H_2O(l) \rightleftharpoons F^-(aq) + H_3O^+(aq)

Reaction B: $BF_3 + F^- \rightarrow BF_4^-$

Reaction C: $NaOH(aq) \rightarrow Na^+(aq) + OH^-(aq)$

For each reaction, the acid-base theory required is:

Reaction A: Brønsted-Lowry (proton transfer from $HF$ to $H_2O$ )
Reaction B: Lewis ( $BF_3$ accepts electron pair from $F^-$ )
Reaction C: Arrhenius ( $NaOH$ produces $OH^-$ in water)

🔑 When classifying, always use the simplest theory that explains the observation.

Problem 1 Practice 🎯

Reaction A: $HF(aq) + H_2O(l) \rightleftharpoons F^-(aq) + H_3O^+(aq)$

Reaction B: $BF_3 + F^- \rightarrow BF_4^-$

🔢 Problem 2: Multi-Step pH Calculation

A chemist prepares the following solutions:

Solution A: 0.035 M $HCl$
Solution B: 0.035 M $NaOH$
Solution C: 50.0 mL of Solution A mixed with 30.0 mL of Solution B

Solution A pH

$[H^+] = 0.035$ M → $pH = -\log(0.035) = 1.46$

Solution B pH

$[OH^-] = 0.035$ M → $pOH = -\log(0.035) = 1.46$ →

Solution C pH

Moles $H^+$ = $(0.050)(0.035) = 1.75 \times 10^{-3}$ mol

Moles $OH^-$ = $(0.030)(0.035) = 1.05 \times 10^{-3}$ mol

Excess $H^+$ = $1.75 \times 10^{-3} - 1.05 \times 10^{-3} = 7.0 \times 10^{-4}$ mol

$[H^+] = \frac{7.0 \times 10^{-4}}{0.080} = 8.75 \times 10^{-3}$ M

$pH = -\log(8.75 \times 10^{-3}) = 2.06$

Problem 2 Practice 🧮

A student mixes 25.0 mL of 0.080 M $HNO_3$ with 15.0 mL of 0.080 M $KOH$ .

1) How many moles of excess $H^+$ remain? (Enter in scientific notation, e.g. 3.5e-3)

2) What is the total volume in liters? (3 decimal places)

3) What is the pH of the resulting solution? (2 decimal places)

📌 Problem 3: Conceptual Reasoning

The pH of Very Dilute Strong Acids

When a strong acid is extremely dilute (e.g., $10^{-8}$ M $HCl$ ), you cannot simply say $pH = 8$ .

⚠️ An acid solution can never have $pH > 7$ !

The autoionization of water contributes $[H^+] = 10^{-7}$ M, which is much larger than the acid's contribution.

Correct approach:

$\boxed{[H^+]_{\text{total}} = [H^+]_{\text{acid}} + [H^+]_{\text{water}} \approx 10^{-8} + 10^{-7} = 1.1 \times 10^{-7} \text{ M}}$

$pH = -\log(1.1 \times 10^{-7}) = 6.96$

💡 This is slightly below 7, as expected for an acidic solution.

Conceptual Check 🎯

Workshop Synthesis 🔍

🔑 Why this matters: This review mirrors the AP exam format — expect questions that require you to connect theory, calculations, and conceptual reasoning in a single problem.

What You'll Master in Part 7

Tackling AP-style multiple choice across all acid-base topics
Writing free-response explanations using proper chemistry terminology
Identifying common AP traps and avoiding them

📋 Complete Summary

Three Acid-Base Theories

Theory	Acid	Base	Scope
Arrhenius	Produces $H^+$	Produces $OH^-$	Aqueous only
Brønsted-Lowry	Proton donor	Proton acceptor	Any solvent
Lewis	e⁻ pair acceptor	e⁻ pair donor	Broadest

Key Equations

$\boxed{pH = -\log[H^+] \qquad pOH = -\log[OH^-]}$

$\boxed{pH + pOH = 14 \qquad K_w = [H^+][OH^-] = 1.0 \times 10^{-14}}$

$[H^+] = 10^{-pH} \qquad [OH^-] = 10^{-pOH}$

Strong Acid/Base Rules

Strong acids: $HCl, HBr, HI, HNO_3, H_2SO_4, HClO_4$

🔑 Strong = complete dissociation. No $K_a$ or $K_b$ needed — just use the concentration directly.

AP-Style Questions — Set 1 🎯

AP Calculation Practice 🧮

1) What is the pH of a solution made by mixing 100.0 mL of 0.15 M $HCl$ with 75.0 mL of 0.15 M $NaOH$ ? (2 decimal places)

2) A solution has a pH of 11.50. What is $[H^+]$ ? (Enter in scientific notation, e.g. 4.7e-9)

3) What volume (mL) of 0.20 M $NaOH$ is needed to exactly neutralize 50.0 mL of 0.10 M $H_2SO_4$ ? (Enter as whole number)

AP-Style Questions — Set 2 🎯

Comprehensive Review 🔍

Final Exit Quiz — Acid-Base Theories & pH ✅

Molecules with lone pairs on N, O, S, or halide ions

[F e (CN)_{6}]^{3 -}

pK_w = -\log K_w = 14

pH = 14 - 2.00 = 12.00

pH = 14 - 2.00 = 12.00

pH = 14 - 1.46 = 12.54

Strong bases: Group 1 hydroxides +

Ca(OH)_2, Sr(OH)_2, Ba(OH)_2

[H^+] = C_{acid}

for monoprotic strong acids

[OH^-] = nC_{base}

where

n

= number of

OH^-

per formula unit

Acid-Base Theories and pH Scale

Acid-Base Theories and pH Scale - Complete Interactive Lesson

Part 1: Arrhenius & Brønsted-Lowry

🧪 Arrhenius Acids and Bases

Three Acid-Base Theories — Where We're Headed

What You'll Master in Part 1

📖 The Arrhenius Definition

🧪 Common Arrhenius Acids

Strong Acids (Complete Dissociation)

⚛️ The Hydronium Ion

📌 Arrhenius Neutralization

Examples

Part 2: Conjugate Acid-Base Pairs

🔄 Brønsted-Lowry Acids and Bases

Arrhenius vs. Brønsted-Lowry

Part 3: The pH Scale

🔬 Lewis Acids and Bases

How the Three Theories Compare

Part 4: Strong Acids & Bases

📊 The pH Scale

The pH Scale at a Glance

Part 5: Calculating pH & pOH

💪 Strong Acids and Bases — pH Calculations

Strong = Complete Dissociation

Part 6: Problem-Solving Workshop

🛠️ Problem-Solving Workshop

Problem Types You'll Practice

What You'll Master in Part 6

🧪 Problem 1: Identifying Acid-Base Behavior

Part 7: Synthesis & AP Review

🎓 Synthesis & AP Review

Everything Comes Together

Key Features

Limitations

Common Strong Bases

Autoionization of Water

Double Replacement Pattern

What You'll Master in Part 2

📖 The Brønsted-Lowry Definition

Key Advantage

Example: HClHClHCl in Water

Example: NH3NH_3NH3​ in Water

🧪 Conjugate Acid-Base Pairs

Examples

Critical Rule

Strength Relationship

⚗️ Identifying Conjugate Pairs in Reactions

Steps to Identify

What You'll Master in Part 3

📖 The Lewis Definition

Comparison of All Three Theories

Key Insight

🧪 Common Lewis Acids

1. Metal Cations

2. Molecules with Incomplete Octets

3. Protons (H+H^+H+)

✏️ Common Lewis Bases

🔗 Coordinate Covalent Bonds

In Coordination Chemistry

What You'll Master in Part 4

🔗 pH, pOH, and Their Relationship

pH Definition

pOH Definition

The Key Relationship

Interpreting pH

🔢 pH Calculations

From [H+][H^+][H+] to pH

From pH to [H+][H^+][H+]

From [OH−][OH^-][OH−] to pH

The "p" Notation

📌 Significant Figures in pH

Examples

What You'll Master in Part 5

🧪 pH of Strong Acids

Example 1

Example 2

Diprotic Strong Acid (H2SO4H_2SO_4H2​SO4​)

📌 pH of Strong Bases

Example 1

Group 2 Hydroxides

Example: $HCl$ in Water

Example: $NH_3$ in Water

3. Protons ( $H^+$ )

From $[H^+]$ to pH

From pH to $[H^+]$

From $[OH^-]$ to pH

Diprotic Strong Acid ( $H_2SO_4$ )