Chemistry Study Guide
Acid-Base
Chemistry
Brønsted-Lowry theory, conjugate pairs, pH & pOH, titration
Core Concepts
Brønsted-Lowry Theory
Acids are proton donors; bases are proton acceptors. Every acid-base reaction is a proton (H⁺) transfer.
Conjugate Pairs
Acid loses H⁺ → becomes its conjugate base. Base gains H⁺ → becomes its conjugate acid.
Strong vs Weak Acids
Strong acids dissociate 100% — single arrow (→). Weak acids partially dissociate — double arrow (⇌).
pH & pOH
pH = −log[H⁺] · pOH = −log[OH⁻] · pH + pOH = 14 at 25°C.
Titration
moles = M × V(L). Use mole ratios from the balanced equation to find moles of analyte.
Neutralization
Acid + Base → Salt + Water. The equivalence point is where moles acid = moles base exactly.
Flashcards
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Term
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Definition
All Terms
Key Equations
H⁺ ≡ H₃O⁺ · All equations will be provided on the check-in
pH = −log[H⁺]
Find pH when you know [H⁺]
★ Core
[H⁺] = 10−pH
Find [H⁺] when you know pH
★ Core
pOH = −log[OH⁻]
Find pOH when you know [OH⁻]
Parallel
[OH⁻] = 10−pOH
Find [OH⁻] when you know pOH
Parallel
pH + pOH = 14
Find one if you have the other (25°C)
★ Core
moles = M × V(L)
Calculate moles of titrant
Titration
Calculator Tip
⚠️
Use the negative (±) key at the bottom of your calculator — not the subtraction key on the right side. To enter 10−5: press 10 ^ (±) 5. They are NOT the same key!
Strong Acids — Memorize All 6
HCl
Hydrochloric
HBr
Hydrobromic
HI
Hydroiodic
HNO₃
Nitric
H₂SO₄
Sulfuric
HClO₄
Perchloric
The pH Scale
Lower pH = more acidic = higher [H⁺] · Each unit = 10× concentration change
AcidicpH < 7
[H⁺] > [OH⁻]
[H⁺] > [OH⁻]
NeutralpH = 7
[H⁺] = [OH⁻]
[H⁺] = [OH⁻]
BasicpH > 7
[OH⁻] > [H⁺]
[OH⁻] > [H⁺]
Example Solutions
Key Relationships
pH decreases →
[H⁺] increases — solution becomes more acidic. Each 1-unit drop = 10× more H⁺.
pH increases →
[OH⁻] increases — solution becomes more basic. pOH decreases as pH increases.
pH Calculator
Enter any one value and get all four instantly.
pH
—
[H⁺] (M)
—
pOH
—
[OH⁻] (M)
—
Practice Problems
Based directly on your Check In #4 practice set.
0
of 0 answered
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In-Depth Review
Choose a topic for a full deep-dive: concepts, examples, and worked problems.
01
Brønsted-Lowry Theory
Acids and bases defined by proton (H⁺) donation and acceptance — not just by producing H⁺ or OH⁻ in water.
Core Definitions
Brønsted-Lowry Acid
A species that donates a proton (H⁺) to another species. The acid loses H⁺ in the reaction. Written with a single (→) or double (⇌) arrow depending on strength.
Brønsted-Lowry Base
A species that accepts a proton (H⁺) from another species. The base gains H⁺ in the reaction. Every proton transfer requires both an acid and a base simultaneously.
H⁺ vs H₃O⁺ — They're the Same Thing
A bare proton (H⁺) doesn't exist alone in water — it immediately bonds to a water molecule to form hydronium (H₃O⁺).
H⁺ + H₂O → H₃O⁺
So when you write [H⁺] and [H₃O⁺], they mean the exact same concentration. Both notations are used interchangeably — your equations sheet confirms this: H⁺ ≡ H₃O⁺.
H⁺ + H₂O → H₃O⁺
So when you write [H⁺] and [H₃O⁺], they mean the exact same concentration. Both notations are used interchangeably — your equations sheet confirms this: H⁺ ≡ H₃O⁺.
Annotated Example Reactions
Strong Acid + Water — single arrow, goes to completion
HCl + H₂O → H₃O⁺ + Cl⁻
HCl = acid (donates H⁺)
H₂O = base (accepts H⁺)
H₃O⁺ = conjugate acid
Cl⁻ = conjugate base
Weak Acid + Water — double arrow, equilibrium established
CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
CH₃COOH = acid (donor)
H₂O = base (acceptor)
Most molecules do not dissociate — equilibrium favors reactants
Water is Amphoteric
Amphoteric
Water can act as either an acid or a base depending on what it's reacting with. This makes it amphoteric. A neutral pH does NOT mean a substance can't donate H⁺ — water does so constantly through autoionization.
Water as BASE (accepts H⁺ from HCl):
HCl + H₂O → H₃O⁺ + Cl⁻
HCl + H₂O → H₃O⁺ + Cl⁻
Water as ACID (donates H⁺ to NH₃):
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
Autoionization of water:
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
One water acts as acid, one as base. At 25°C this gives [H⁺] = [OH⁻] = 1×10⁻⁷ M, so pH = 7.
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
One water acts as acid, one as base. At 25°C this gives [H⁺] = [OH⁻] = 1×10⁻⁷ M, so pH = 7.
Key Takeaways
- The acid donates H⁺; the base accepts H⁺ — every reaction is a proton transfer.
- H⁺ ≡ H₃O⁺ — always interchangeable in equations and calculations.
- Strong acids use → (100% dissociation); weak acids use ⇌ (partial).
- Water is amphoteric — acts as acid or base depending on context. Neutral pH ≠ unable to donate H⁺.
02
Conjugate Acid-Base Pairs
Every B-L reaction produces two conjugate pairs. They always differ by exactly one H⁺.
The Golden Rule
Conjugate Base
Formed when an acid loses H⁺. Take the acid, remove one H⁺, and adjust the charge by −1.
HF → F⁻ (lost H⁺, charge −1)
H₂O → OH⁻ (lost H⁺, charge −1)
HF → F⁻ (lost H⁺, charge −1)
H₂O → OH⁻ (lost H⁺, charge −1)
Conjugate Acid
Formed when a base gains H⁺. Take the base, add one H⁺, and adjust the charge by +1.
NH₃ → NH₄⁺ (gained H⁺, charge +1)
CN⁻ → HCN (gained H⁺, charge 0)
NH₃ → NH₄⁺ (gained H⁺, charge +1)
CN⁻ → HCN (gained H⁺, charge 0)
Step-by-Step: How to Identify All Four Species
1
Find the proton donor (the ACID)
Look at the reactants. Which species loses an H⁺ to become a product? That's the B-L acid. Compare left-side and right-side species — the acid has one more H than its conjugate base.
2
Find the proton acceptor (the BASE)
The other reactant is the B-L base — it gains H⁺. Look at the other product to confirm: it should have one more H than the base (that's the conjugate acid).
3
Write Conjugate Base = Acid − H⁺
Remove one H from the acid formula and subtract 1 from the charge.
HCO₃⁻ − H⁺ = CO₃²⁻ (charge goes from −1 to −2)
4
Write Conjugate Acid = Base + H⁺
Add one H to the base formula and add 1 to the charge.
OH⁻ + H⁺ = H₂O (charge goes from −1 to 0)
All Four Species — Visual Layout
HF + CN⁻ ⇌ F⁻ + HCN
Pair 1 (differ by H⁺)
HF
−H⁺ →
F⁻
Acid → Conjugate Base
Pair 2 (differ by H⁺)
CN⁻
+H⁺ →
HCN
Base → Conjugate Acid
Full Example Table
| Reaction | B-L Acid | B-L Base | Conj. Acid | Conj. Base |
|---|---|---|---|---|
| NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ | H₂O | NH₃ | NH₄⁺ | OH⁻ |
| HF + CN⁻ ⇌ F⁻ + HCN | HF | CN⁻ | HCN | F⁻ |
| HCO₃⁻ + OH⁻ ⇌ CO₃²⁻ + H₂O | HCO₃⁻ | OH⁻ | H₂O | CO₃²⁻ |
| H₂PO₄⁻ + H₂O ⇌ HPO₄²⁻ + H₃O⁺ | H₂PO₄⁻ | H₂O | H₃O⁺ | HPO₄²⁻ |
Completing a Missing Species
WORKED EXAMPLECO₃²⁻ + H₂O ⇌ HCO₃⁻ + ?
Find: the missing product.
1
H₂O is the acid (it donates H⁺ to CO₃²⁻). CO₃²⁻ is the base.
2
Conjugate base of H₂O = H₂O − H⁺ = OH⁻
3
Check: CO₃²⁻ + H⁺ = HCO₃⁻ ✓ (that's the conjugate acid, shown on the right)
Missing species = OH⁻ (conjugate base of H₂O)
Key Takeaways
- Conjugate pairs always differ by exactly one H⁺ — check charges too.
- Acid − H⁺ → Conjugate Base (charge decreases by 1).
- Base + H⁺ → Conjugate Acid (charge increases by 1).
- Every B-L reaction has exactly two conjugate pairs.
- The "stronger" the acid, the weaker its conjugate base (and vice versa).
03
Strong vs Weak Acids
The difference is the extent of dissociation — how completely the acid transfers H⁺ to water.
Head-to-Head Comparison
| Property | Strong Acid (HCl) | Weak Acid (CH₃COOH) |
|---|---|---|
| Dissociation | 100% complete | Partial (<5% typically) |
| Arrow used | → (single) | ⇌ (double) |
| [H⁺] from 0.1 M | 0.1 M | ≈ 0.0013 M |
| pH of 0.1 M | 1.0 | ≈ 2.87 |
| More acidic? | Yes — lower pH | No |
| Equilibrium? | No — reaction is complete | Yes — most stays as HA |
Dissociation Equations
Strong acid — single arrow, 100% dissociation
HCl + H₂O → H₃O⁺ + Cl⁻
0.1 M HCl produces exactly 0.1 M H₃O⁺ → pH = 1.0
Weak acid — double arrow, equilibrium favors reactants
CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻
0.1 M CH₃COOH produces only ≈0.0013 M H₃O⁺ → pH ≈ 2.87
The 6 Strong Acids — Know All of Them
Hydrochloric
The most common strong acid. Fully ionizes: HCl → H⁺ + Cl⁻
Hydrobromic
Same behavior as HCl — complete dissociation in water.
Hydroiodic
Stronger than HCl and HBr due to the larger iodide ion.
Nitric
Strong acid used in synthesis and as an oxidizing agent.
Sulfuric
Diprotic — the first ionization is strong. Second is weak.
Perchloric
Actually the strongest of the six. Complete dissociation.
💡 Everything else is a weak acid. If you don't recognize it as one of the six above, treat it as weak (partial dissociation, double arrow, [H⁺] < [acid]).
Why This Matters for pH
WORKED EXAMPLECompare pH of two 0.1 M solutions
Given: 0.1 M HCl and 0.1 M acetic acid (CH₃COOH). Which has the lower pH?
1
HCl is strong → dissociates 100% → [H⁺] = 0.1 M → pH = −log(0.1) = 1.0
2
CH₃COOH is weak → only ~1.3% dissociates → [H⁺] ≈ 0.0013 M → pH = −log(0.0013) ≈ 2.87
3
Lower pH = more acidic. HCl is more acidic.
HCl has the lower pH (1.0 vs 2.87) — same concentration, but more H⁺ because it dissociates completely.
Key Takeaways
- Strong acids dissociate 100% — use → and assume [H⁺] = [acid].
- Weak acids dissociate partially — use ⇌ and know [H⁺] << [acid].
- At the same concentration, a strong acid always has a lower pH than a weak acid.
- Memorize all 6 strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄. Everything else is weak.
- The arrow type (→ vs ⇌) communicates how far the reaction proceeds.
04
pH, pOH & Calculations
A logarithmic scale that compresses an enormous range of H⁺ concentrations (10⁰ to 10⁻¹⁴) into a 0–14 number line.
All Four Formulas
pH = −log[H⁺]Use when you know [H⁺] and want pH
[H⁺] = 10−pHUse when you know pH and want [H⁺]
pOH = −log[OH⁻]Use when you know [OH⁻] and want pOH
[OH⁻] = 10−pOHUse when you know pOH and want [OH⁻]
pH + pOH = 14Always true at 25°C — find one if you have the other
Why Logarithms?
H⁺ concentrations in real solutions span 14 orders of magnitude — from 1 M (very strong acid) to 10⁻¹⁴ M (very strong base). Comparing 0.000001 M vs 0.0000001 M is hard. But pH 6 vs pH 7 is instantly readable.
Each 1-unit increase in pH = 10× decrease in [H⁺]. pH 4 is 10× more acidic than pH 5, and 100× more acidic than pH 6.
Each 1-unit increase in pH = 10× decrease in [H⁺]. pH 4 is 10× more acidic than pH 5, and 100× more acidic than pH 6.
Interpreting pH Values
Acidic (pH < 7)
[H⁺] > [OH⁻]
More H⁺ ions than OH⁻ ions. Lower pH = higher [H⁺] = more acidic.
More H⁺ ions than OH⁻ ions. Lower pH = higher [H⁺] = more acidic.
Neutral (pH = 7)
[H⁺] = [OH⁻] = 10⁻⁷ M
Exactly balanced. Pure water at 25°C. pH = pOH = 7.
Exactly balanced. Pure water at 25°C. pH = pOH = 7.
Basic (pH > 7)
[OH⁻] > [H⁺]
More OH⁻ ions than H⁺ ions. Higher pH = higher [OH⁻] = more basic.
More OH⁻ ions than H⁺ ions. Higher pH = higher [OH⁻] = more basic.
Step-by-Step Calculation Workflow
A
pH → [H⁺]
Use [H⁺] = 10−pH. On your calculator: enter the pH, press +/− (negation key, not minus), then press 10x or use the ^ key.
pH = 5 → [H⁺] = 10⁻⁵ = 1×10⁻⁵ M
B
[H⁺] → pH
Use pH = −log[H⁺]. Take the log of [H⁺], then negate it.
[H⁺] = 1×10⁻⁵ M → pH = −log(10⁻⁵) = −(−5) = 5
C
pH → pOH → [OH⁻]
Use pOH = 14 − pH, then [OH⁻] = 10−pOH.
pH = 5 → pOH = 9 → [OH⁻] = 10⁻⁹ M
Full Worked Examples
EXAMPLE 1Black coffee — pH = 5.0
Find: [H⁺], pOH, and [OH⁻].
1
[H⁺] = 10⁻⁵·⁰ = 1×10⁻⁵ M
2
pOH = 14 − 5.0 = 9.0
3
[OH⁻] = 10⁻⁹·⁰ = 1×10⁻⁹ M
4
pH 5 < 7 → acidic solution ✓
[H⁺] = 1×10⁻⁵ M · pOH = 9 · [OH⁻] = 1×10⁻⁹ M
EXAMPLE 2Lemon juice — pH = 2.4
Find: [H⁺], pOH, and [OH⁻].
1
[H⁺] = 10⁻²·⁴ ≈ 3.98×10⁻³ M (use calculator: 10^(±2.4))
2
pOH = 14 − 2.4 = 11.6
3
[OH⁻] = 10⁻¹¹·⁶ ≈ 2.51×10⁻¹² M
[H⁺] ≈ 3.98×10⁻³ M · pOH = 11.6 · [OH⁻] ≈ 2.5×10⁻¹² M
⚠️
Calculator Warning: When entering a negative exponent (like 10−5), use the +/− negation key at the bottom of your calculator — NOT the subtraction key on the right side. They look similar but behave completely differently. The subtraction key will give you a wrong answer every time.
Key Takeaways
- pH + pOH = 14 always (at 25°C) — know one, find the other instantly.
- Lower pH = more H⁺ = more acidic. Each unit = 10× change in [H⁺].
- Use 10−pH to get [H⁺]; use −log[H⁺] to get pH.
- pH < 7 = acidic, pH = 7 = neutral, pH > 7 = basic.
- On your calculator: always use the negation key (+/−), never the subtraction key.
05
Neutralization & Titration
When acids and bases react completely, they cancel each other out. Titration is how we use this to measure an unknown concentration.
Neutralization
Neutralization Reaction
An acid and a base react to form a salt + water. The H⁺ from the acid combines with the OH⁻ from the base to make H₂O.
Pattern: Acid + Base → Salt + Water
Pattern: Acid + Base → Salt + Water
Equivalence Point
The point where moles of acid exactly equal moles of base (for a 1:1 reaction). The solution is not necessarily pH 7 — that depends on the salt formed.
Example neutralization reactions
HCl + NaOH → NaCl + H₂O
H₂SO₄ + 2 NaOH → Na₂SO₄ + 2 H₂O
H₂SO₄ needs 2 moles of NaOH — always use mole ratios from the balanced equation!
Titration Calculation — Step by Step
1
Calculate moles of titrant (known solution)
The titrant is the solution with known molarity (M) added from a buret.
moles = M × V(L) ← must convert mL to L first!
2
Apply the mole ratio from the balanced equation
For 1:1 reactions (HCl + NaOH), moles analyte = moles titrant. For other ratios, multiply accordingly.
H₂SO₄ + 2 NaOH → moles H₂SO₄ = moles NaOH ÷ 2
3
Convert moles to grams (if asked)
Multiply moles of analyte by its molar mass.
grams = moles × molar mass (g/mol)
Full Worked Example
WORKED EXAMPLE25.0 mL of 0.200 M NaOH neutralizes HCl
Find: moles of NaOH used, and moles of HCl neutralized.
1
Convert volume: 25.0 mL = 0.0250 L
2
Moles of NaOH: 0.200 M × 0.0250 L = 0.00500 mol
3
Balanced equation: HCl + NaOH → NaCl + H₂O → 1:1 mole ratio
4
Moles HCl = moles NaOH = 0.00500 mol
5
If asked for grams of HCl: 0.00500 mol × 36.46 g/mol = 0.182 g
0.00500 mol NaOH used → 0.00500 mol HCl neutralized → 0.182 g HCl
The Role of an Indicator
Indicator
A weak acid (or base) that has different colors in its protonated vs deprotonated form. Added in tiny amounts to a titration so it doesn't significantly affect the pH. It changes color near the equivalence point, letting you visually see when neutralization is complete — even though you can't see H⁺ directly.
Why does the color change? As you add base, [H⁺] drops. When [H⁺] falls below a threshold, the indicator (a weak acid) loses its H⁺ and changes form — and the two forms absorb different wavelengths of light, giving different colors. The transition zone is usually within ±1 pH unit of the equivalence point.
Key Takeaways
- Neutralization: Acid + Base → Salt + Water. The H⁺ and OH⁻ combine to form H₂O.
- Equivalence point = moles acid = moles base (for 1:1); use the balanced equation's ratio for other reactions.
- Key formula: moles = M × V(L) — always convert mL → L first.
- To find grams: moles × molar mass.
- Indicators are weak acids that change color near the equivalence point to signal titration completion.
How to Tell if Something is an Acid or Base
The Core Question
Ask: does this substance donate H⁺ or accept H⁺? If it gives away a proton → acid. If it grabs a proton → base. Everything else follows from this single test.
Strategy 1 — Look at the Formula
1
Starts with H (and is not water)? Almost certainly an acid. It has a proton it can donate. Examples: HCl, HNO₃, H₂SO₄, CH₃COOH.
2
Ends with OH? Likely a base. The OH⁻ group will accept H⁺ and form water. Examples: NaOH, KOH, Ca(OH)₂.
3
Contains NH₃ or NH₂? Base. The nitrogen lone pair grabs H⁺. Examples: NH₃, methylamine (CH₃NH₂).
4
Carbonate or bicarbonate (CO₃²⁻, HCO₃⁻)? Base — these ions accept protons.
Strategy 2 — Check the Reaction Context
In a Brønsted-Lowry reaction, roles are assigned by what happens, not just by formula. Watch the arrow:
HF + H₂O ⇌ H₃O⁺ + F⁻
HF = acid (gives H⁺)
H₂O = base (takes H⁺)
H₃O⁺ = conjugate acid
F⁻ = conjugate base
Strategy 3 — Use pH (if you have it)
pH < 7
Acidic solution
pH = 7
Neutral (pure water)
pH > 7
Basic solution
Amphoteric Species — Can Be Either
Amphoteric (Amphiprotic)
A species that can both donate and accept H⁺ depending on what it reacts with. Water (H₂O) and the bicarbonate ion (HCO₃⁻) are the classic examples.
H₂O acting as acid: H₂O + NH₃ → OH⁻ + NH₄⁺
H₂O acting as base: HCl + H₂O → H₃O⁺ + Cl⁻
Quick-Reference Table
| Clue | Likely Role | Example |
|---|---|---|
| Formula starts with H (not H₂O) | Acid | HCl, H₂SO₄, HNO₃ |
| Formula ends with OH | Base | NaOH, KOH, Mg(OH)₂ |
| Contains NH₃ / NH₂ group | Base | NH₃, CH₃NH₂ |
| Carbonate / bicarbonate ion | Base | Na₂CO₃, NaHCO₃ |
| pH < 7 measured | Acid | Lemon juice (pH ≈ 2) |
| pH > 7 measured | Base | Bleach (pH ≈ 13) |
| Donates H⁺ in reaction | Acid | Any B-L acid |
| Accepts H⁺ in reaction | Base | Any B-L base |
Worked Example — Classify Each Species
In the reaction: NH₄⁺ + OH⁻ → NH₃ + H₂O, identify the acid and the base.
1Ask: who gives H⁺? NH₄⁺ loses one H⁺ to become NH₃. So NH₄⁺ is the acid.
2Ask: who takes H⁺? OH⁻ gains H⁺ and becomes H₂O. So OH⁻ is the base.
3Conjugate pairs: NH₄⁺ / NH₃ and OH⁻ / H₂O.
Acid = NH₄⁺ | Base = OH⁻
Key Takeaways
- H⁺ donor = acid; H⁺ acceptor = base — this is the only rule you truly need.
- Formula clues (starts with H, ends with OH) are shortcuts, not absolutes — always confirm in context.
- pH < 7 → acidic; pH > 7 → basic; pH = 7 → neutral at 25 °C.
- Amphoteric species (H₂O, HCO₃⁻) can act as either — context decides.
Visualize — See the Concepts
Interactive pH Scale
Hover any marker to see the substance's details. The color gradient mirrors real litmus paper.
← More Acidic [H⁺] increases
Neutral
More Basic → [OH⁻] increases
Strong vs Weak: Dissociation Visual
Each circle is one molecule. Red = H⁺ ions, Grey = intact molecule. Both beakers are 0.1 M.
HCl (Strong Acid)
100% dissociated → pH = 1.0
CH₃COOH (Weak Acid)
~1.3% dissociated → pH ≈ 2.87
Proton Transfer in Action
Step through a Brønsted-Lowry reaction one move at a time.
HFH⁺
+
CN⁻
⇌
HCN
+
F⁻
B-L Acid (HF)
B-L Base (CN⁻)
Conjugate Acid
Conjugate Base
Click "Show proton transfer" to watch H⁺ move from HF to CN⁻.
Formula Map
[H⁺]
measured / known
−log ↓ 10^(−x) ↑
pH
= −log[H⁺]
+ pOH = 14
pOH
= −log[OH⁻]
−log ↓ 10^(−x) ↑
[OH⁻]
measured / known
What to take away from the visuals
- The pH scale is logarithmic — each unit is a 10× change in [H⁺].
- Strong acids produce vastly more H⁺ than weak acids at the same concentration.
- Proton transfer is a single H⁺ moving — everything else about the species stays the same.
- All four variables ([H⁺], pH, pOH, [OH⁻]) are linked — know one, find all four.
Titration Curve — Strong Acid + Strong Base
Setup
25.0 mL of 0.100 M HCl titrated with 0.100 M NaOH. Drag the slider to add base and watch the pH change.
Volume NaOH0.0 mL
pH1.00
pOH13.00
RegionBefore equivalence
Reading the Curve
1
Before equivalence (0–25 mL): excess H⁺ controls pH. Curve rises slowly as HCl is neutralized.
2
Equivalence point (25.0 mL): moles NaOH = moles HCl. For strong/strong: pH = 7. The steep vertical jump happens here.
3
After equivalence (25–50 mL): excess OH⁻ controls pH. Curve flattens as the base is diluted.
4
Indicator choice: pick an indicator whose color-change range (pKa ± 1) overlaps the steep jump. Phenolphthalein (pH 8.2–10) works well here.
Key Takeaways
- The equivalence point is where moles acid = moles base — for a 1:1 reaction, use n = MV.
- Strong acid + strong base → salt + water. At equivalence: pH = 7.
- The steep vertical region is where the indicator changes color — that's the titration endpoint.
- A weak acid + strong base curve shifts left and equivalence pH > 7 (the conjugate base makes it basic).
Cheat Sheet
Everything you need on one page — formulas, rules, strong acids, and quick classification.
Key Formulas
pH = −log[H⁺]
[H⁺] = 10−pH
pOH = −log[OH⁻]
[OH⁻] = 10−pOH
pH + pOH = 14
at 25 °C
moles = M × V(L)
titration — convert mL first!
pH Quick Guide
0 – 6
Acidic
[H⁺] > [OH⁻]
7
Neutral
[H⁺] = [OH⁻]
8 – 14
Basic
[OH⁻] > [H⁺]
The 6 Strong Acids
HCl
HBr
HI
HNO₃
H₂SO₄
HClO₄
Everything else is a weak acid (partial dissociation, ⇌).
Common Strong Bases
NaOH
KOH
Ca(OH)₂
Ba(OH)₂
Brønsted-Lowry Rules
Aciddonates H⁺ to another species
Baseaccepts H⁺ from another species
Conj. Acidbase + H⁺ (formed in reaction)
Conj. Baseacid − H⁺ (formed in reaction)
Conjugate pairs differ by exactly one H⁺.
Amphoteric
H₂O and HCO₃⁻ can act as either acid or base depending on context.
Reaction Patterns
Neutralization
Acid + Base → Salt + H₂O
Strong acid
HA → H⁺ + A⁻ (single arrow, 100%)
Weak acid
HA ⇌ H⁺ + A⁻ (double arrow, partial)
Autoionization
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
Identify Acid or Base — Quick Clues
Starts with H (not H₂O)Likely AcidHCl, H₂SO₄
Ends in OHLikely BaseNaOH, KOH
Contains NH₃/NH₂Likely BaseNH₃
Carbonate / bicarbonateLikely BaseNa₂CO₃
pH < 7Acidiclemon juice
pH > 7Basicbleach
Donates H⁺ in reactionB-L Acidcontext-based
Accepts H⁺ in reactionB-L Basecontext-based
Common Mistakes to Avoid
✗Forgetting to convert mL → L before using moles = M × V.
✗Using pH instead of pOH (or vice versa) when solving for concentration.
✗Assuming "neutral pH" means water can't donate H⁺ — it's amphoteric.
✗Treating weak acids as if they fully dissociate — they don't (use ⇌, not →).
✗Confusing conjugate acid (base + H⁺) with conjugate base (acid − H⁺).