Acids, Bases, and Salts

Subject: Chemistry
Topic: 4
Cambridge Code: 0620 / 0971 / 5070

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Acids and Bases

Acid Definition

Arrhenius definition:

• Produces H⁺ ions (hydrogen ions)
• In aqueous solution
• Taste sour
• Turn blue litmus red

Lowry-Brønsted definition:

• Proton (H⁺) donor

Base Definition

Arrhenius definition:

• Produces OH⁻ ions (hydroxide ions)
• In aqueous solution
• Taste bitter
• Slippery
• Turn red litmus blue

Lowry-Brønsted definition:

• Proton (H⁺) acceptor

Common Acids

Strong acids:

• Completely ionize
• HCl, HBr, HI (hydrohalic)
• HNO₃ (nitric)
• H₂SO₄ (sulfuric)

Weak acids:

• Partially ionize
• CH₃COOH (acetic)
• HCO₂H (formic)
• H₂CO₃ (carbonic)

Common Bases

Strong bases:

• Group 1 hydroxides (NaOH, KOH)
• Soluble and ionize completely

Weak bases:

• NH₃ (ammonia)
• Amines
• Partially ionize

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pH Scale

pH - Measure of acidity/basicity

$pH = -\log[H^+]$

Scale:

• pH < 7: Acidic
• pH = 7: Neutral (at 25°C)
• pH > 7: Basic (alkaline)

pH Values

Solution	pH
Stomach acid	2
Lemon juice	2.5
Vinegar	3
Tomato juice	4
Coffee	5
Pure water	7
Sea water	8.2
Ammonia	11
Sodium hydroxide	14

Calculating pH

Example: [H⁺] = 10⁻³

• pH = -log(10⁻³) = 3 (acidic)

pOH and Relationship

$pH + pOH = 14$

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Indicators

Indicator - Dye that color changes with pH

Common Indicators

Litmus:

• Red: Acidic
• Blue: Alkaline

Methyl orange:

• Red: pH < 3.1
• Orange: pH 3.1-4.4
• Yellow: pH > 4.4

Phenolphthalein:

• Colorless: pH < 8.2
• Pink: pH 8.2-10
• Purple: pH > 10

Universal indicator:

• Range of colors
• Estimates exact pH

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Neutralization Reactions

Neutralization - Acid + Base → Salt + Water

$HA + BOH → BA + H_2O$

Ionic equation: $H^+ + OH^- → H_2O$

Salt Formation

Salt - Ionic compound from acid + base

Examples:

• HCl + NaOH → NaCl + H₂O
• H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
• HNO₃ + NH₃ → NH₄NO₃

Soluble and Insoluble Salts

Soluble salts:

• All alkali salts
• Most chlorides, nitrates
• Most sulfates

Insoluble salts:

• Formed by precipitation
• AgCl, BaSO₄, CaCO₃

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Acid-Base Titration

Titration - Analytical technique to find concentration

Equipment

• Burette - Measures volume precisely (0.05 mL)
• Pipette - Measures fixed volume
• Conical flask - Contains sample
• Indicator - Shows endpoint

Procedure

1. Pipette fixed volume of acid into conical flask
2. Add indicator (3-5 drops)
3. Fill burette with alkali (note volume)
4. Add alkali from burette, swirl frequently
5. Stop when color change occurs (endpoint)
6. Note final burette volume
7. Repeat until consistent results (concordant)

Volume Calculation

$V_1 = \text{Initial volume}$ $V_2 = \text{Final volume}$ $\text{Volume used} = V_2 - V_1$

Example: Initial 0.00 mL, Final 25.00 mL

• Volume used = 25.00 mL

Concentration Calculation

$n_A = n_B$ $M_A V_A = M_B V_B$

From stoichiometry if coefficients ≠ 1: $\frac{M_A \times V_A}{n_A} = \frac{M_B \times V_B}{n_B}$

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Buffer Solutions

Buffer - Solution resisting pH change

Components

• Weak acid + salt of its conjugate base
• OR weak base + salt of its conjugate acid

Example: CH₃COOH + CH₃COONa

How it Works

If acid added:

• Conjugate base neutralizes: OH⁻ + H⁺ → H₂O

If base added:

• Weak acid neutralizes: H⁺ + OH⁻ → H₂O

Importance

• Blood pH maintenance (7.35-7.45)
• Laboratory solutions
• Industrial processes
• Food preservation

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Key Points

1. Acids: H⁺ donors, pH < 7
2. Bases: OH⁻ producers or H⁺ acceptors, pH > 7
3. pH scale: 0-14
4. Neutralization: Acid + Base → Salt + Water
5. Indicators: Show pH range
6. Titration: Measures concentration
7. Buffers: Resist pH change

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Practice Questions

1. Classify acids/bases as strong/weak
2. Calculate pH from [H⁺]
3. Predict indicator colors
4. Write neutralization equations
5. Calculate simple titration
6. Explain buffer action

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Revision Tips

• Know pH scale clearly
• Learn strong/weak acids/bases
• Know indicator colors
• Practice titration calculations
• Understand buffer concept
• Learn salt formation
• Practice equation writing