🌟 Module 3: Acidic and Basic Nature – The pH Concept

🎯 Objective:

To understand the pH scale, its scientific basis, and how acidic or basic nature of substances impacts daily life—ranging from the human body to agriculture and environment.

🔬 1. What is pH?

🔹 The term pH stands for "potential of Hydrogen" or "power of Hydrogen".

🔹 It is a measure of hydrogen ion concentration $[H⁺]$ in a solution.

🔹 pH is a logarithmic scale:

$$\text{pH} = -\log_{10}[H^+]$$

• If $[H^+]$ is high → pH is low → Acidic.

• If $[OH^-]$ is high → pH is high → Basic (Alkaline).

📈 2. The pH Scale (0–14)

pH Value	Nature of Substance	Example	Universal Indicator Color
0 – 3	Strongly Acidic	Hydrochloric acid (HCl), gastric juice	Red / Dark Orange
4 – 6	Weakly Acidic	Tomato juice, vinegar	Orange to Yellow
7	Neutral	Pure water	Green
8 – 10	Weakly Basic	Baking soda solution	Blue
11 – 14	Strongly Basic	NaOH, KOH solutions	Purple / Violet

🧠 Concept Insight:

• pH < 7 → Acidic

• pH = 7 → Neutral

• pH > 7 → Basic

The farther from 7, the stronger the acid/base.

🌿 3. Importance of pH in Everyday Life

👤 a. pH in the Human Body

• The human body maintains a blood pH around 7.35–7.45.

• Enzymes, hormones, and cell functions require stable pH.

• Drastic changes in pH can be fatal (acidosis or alkalosis).

🔍 Example:

• Stomach acid has a pH of 1.5–3.5 for digestion (HCl).

• Antacids like milk of magnesia (Mg(OH)₂) are basic to neutralize excess acid.

🌱 b. pH of Soil for Plants

• Soil pH affects nutrient absorption.

• Most crops grow well in pH 6 to 7.5.

• Too acidic soil → Add quicklime (CaO).

• Too basic soil → Add organic matter or sulfur.

🔍 Example:

• Tea grows well in acidic soils (pH < 6).

• Farmers often test soil pH before fertilizing.

🦷 c. Tooth Decay and pH

• Bacteria in the mouth produce acids after eating sugary foods.

• If pH of mouth falls below 5.5, tooth enamel starts dissolving.

• Toothpastes are basic to neutralize this acid and prevent cavities.

☔ d. Acid Rain

• Pollutants like SO₂ and NO₂ dissolve in rainwater to form H₂SO₄ and HNO₃.

• Rainwater becomes acidic (pH < 5.6) → Damages buildings, crops, aquatic life.

🔍 Equation:

$$\text{SO}_2 + \text{H}_2O \rightarrow \text{H}_2\text{SO}_3$$ $$\text{H}_2\text{SO}_3 + [O] \rightarrow \text{H}_2\text{SO}_4$$

🌈 4. Universal Indicator and pH Colors

• A universal indicator is a mixture of several indicators that shows a range of colors for different pH levels.

📊 Color Guide:

pH	Color	Nature
1	Red	Strong acid
4	Orange	Weak acid
7	Green	Neutral
9	Blue	Weak base
13	Purple	Strong base

🧪 5. Demonstration Activity – pH Testing

🎲 Materials:

• Universal indicator paper

• Solutions: Vinegar, soap water, lemon juice, milk, baking soda, cola, shampoo, detergent

• Small containers

🔬 Procedure:

1. Dip the indicator strip into each liquid.

2. Compare the color with the standard pH color chart.

3. Record observations.

📋 Sample Observations:

Substance	pH Range	Indicator Color	Nature
Vinegar	3	Orange-Red	Acidic
Lemon juice	2–3	Red	Acidic
Soap water	9–10	Blue	Basic
Milk	6.5–6.8	Pale Yellow-Green	Slightly acidic
Baking soda	8–9	Blue-Green	Weak base
Cola	2–3	Red	Strong acid
Shampoo	5–6	Yellow	Mild acid

💡 Misconception Alert

• A lower pH does NOT mean a “better acid”—it means a stronger acid.

• Acidic ≠ sour taste only. Some acids are dangerous to taste or touch.

• Not all alkaline solutions feel slippery. Some can be corrosive, like NaOH.

🧠 Quick Revision Capsule

• pH measures acidity/basicity using hydrogen ion concentration.

• Universal indicator helps determine pH through color change.

• Neutralization reactions help in controlling pH in soil, mouth, and body.

• pH control is essential in agriculture, healthcare, and environment.

Module 2: Chemical Properties of Acids and Bases

 

🌟 Module 2: Chemical Properties of Acids and Bases

Chapter: Acids, Bases and Salts – Class 10 Science (CBSE)

🎯 Objective:

To understand the chemical behaviour of acids and bases with various classes of compounds including metals, metal oxides, metal carbonates, and non-metal oxides, through balanced chemical equations and real-life examples.

🔬 1. Reaction of Acids with Metals

🔹 General Reaction:

$$\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas (H₂↑)}$$

✅ Example:

$$\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2↑$$

• Hydrogen gas is evolved – it can be tested by bringing a burning splint near the mouth of the test tube. A ‘pop’ sound confirms hydrogen.

• Salt formed depends on the acid and metal used.

🧠 Concept Insight:

Metals displace hydrogen from acids because they are more electropositive. This is a single displacement reaction.

🧪 2. Reaction of Acids with Metal Carbonates and Bicarbonates

🔹 General Reactions:

$$\text{Acid} + \text{Metal carbonate} \rightarrow \text{Salt} + \text{Carbon dioxide (CO₂)} + \text{Water}$$
$$\text{Acid} + \text{Metal bicarbonate} \rightarrow \text{Salt} + \text{CO}_2 + \text{H}_2O$$

✅ Example 1:

$$\text{Na}_2\text{CO}_3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{CO}_2↑ + \text{H}_2O$$

✅ Example 2:

$$\text{NaHCO}_3 + \text{HCl} \rightarrow \text{NaCl} + \text{CO}_2↑ + \text{H}_2O$$

🔍 CO₂ Detection Test:

Pass the evolved gas through lime water (Ca(OH)₂). If it turns milky, CO₂ is present:

$$\text{Ca(OH)}_2 + \text{CO}_2 \rightarrow \text{CaCO}_3↓ + \text{H}_2O$$

🔥 3. Reaction of Acids with Metal Oxides (Basic Oxides)

Metal oxides are basic in nature and neutralize acids.

🔹 General Reaction:

$$\text{Acid} + \text{Metal oxide} \rightarrow \text{Salt} + \text{Water}$$

✅ Example:

$$2\text{HCl} + \text{CuO} \rightarrow \text{CuCl}_2 + \text{H}_2O$$

🧠 Concept Insight:
This is a neutralisation reaction where metal oxides act like bases.

🌫️ 4. Reaction of Bases with Non-metal Oxides (Acidic Oxides)

Non-metal oxides like CO₂ are acidic in nature and react with bases.

🔹 General Reaction:

$$\text{Base} + \text{Non-metal oxide} \rightarrow \text{Salt} + \text{Water}$$

✅ Example:

$$\text{Ca(OH)}_2 + \text{CO}_2 \rightarrow \text{CaCO}_3 + \text{H}_2O$$

• This reaction also forms insoluble calcium carbonate, making lime water turn milky.

• Indicates that CO₂ behaves as an acidic oxide.

🧪 Activity: Baking Soda and Vinegar Reaction

🎲 Materials:

• Baking soda (sodium bicarbonate – NaHCO₃)

• Vinegar (contains acetic acid – CH₃COOH)

• Transparent glass

• Spoon

🧪 Procedure:

1. Take 1 spoon of baking soda in a glass.

2. Add 2–3 spoons of vinegar.

3. Observe the fizzing and bubble formation.

🔬 Reaction:

$$\text{NaHCO}_3 + \text{CH}_3\text{COOH} \rightarrow \text{CH}_3\text{COONa} + \text{CO}_2↑ + \text{H}_2O$$

💡 Observation:

• CO₂ gas causes fizzing.

• It is the same principle used in baking, where CO₂ causes dough to rise.

📌 Summary of Key Reactions

Reaction Type	Example Equation
Acid + Metal	Zn + HCl → ZnCl₂ + H₂
Acid + Carbonate/Bicarbonate	Na₂CO₃ + HCl → NaCl + CO₂ + H₂O
Acid + Metal Oxide	CuO + HCl → CuCl₂ + H₂O
Base + Non-metal Oxide	Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

🧠 Think Like a Scientist

🔹 Q: Why is CO₂ treated as an acidic oxide?

🔹 A: Because it reacts with bases to form salt and water—just like an acid.

🔹 Q: Why is fizzing seen when vinegar is added to baking soda?

🔹 A: Because CO₂ gas is released due to acid-base neutralisation.

🧠 Misconceptions Cleared

• Not all metal oxides are acidic; metal oxides are generally basic.

• Not all acids react with all metals—unreactive metals like copper do not displace hydrogen from dilute acids like HCl.

• Baking soda is a weak base, not a neutral compound.

Module 1: Introduction to Acids and Bases

 

🔬 Module 1: Introduction to Acids and Bases

(Class 10 CBSE Science – Chapter: Acids, Bases, and Salts)

🌟 1. What are Acids and Bases?

Acids and Bases are two important categories of chemical substances that show distinctly different chemical and physical properties.

📘 Definition:

• Acids: Substances that release hydrogen ions (H⁺) in an aqueous solution.
  ➤ Example: Hydrochloric acid (HCl) dissociates in water as:
  HCl → H⁺ + Cl⁻

• Bases: Substances that release hydroxide ions (OH⁻) in an aqueous solution.
  ➤ Example: Sodium hydroxide (NaOH) dissociates as:
  NaOH → Na⁺ + OH⁻

📌 Note: The presence of H⁺ or OH⁻ ions in water is responsible for the acidic or basic nature of the substance.

🔍 2. Physical Properties

Property	Acids	Bases
Taste	Sour (Don't taste in lab!)	Bitter (Don't taste in lab!)
Touch	Corrosive	Soapy or slippery
Effect on litmus	Turns blue to red	Turns red to blue
Electrical conductivity	Good conductor (due to H⁺ ions)	Good conductor (due to OH⁻ ions)

🧪 3. Common Examples from Daily Life

Acids	Bases
Lemon juice (citric acid)	Soap (sodium hydroxide)
Vinegar (acetic acid)	Baking soda (sodium bicarbonate)
Curd (lactic acid)	Toothpaste (mild base)
Tamarind (tartaric acid)	Limewater (calcium hydroxide)

🧫 4. Indicators: Detecting Acids and Bases

Indicators are substances that change color in acidic or basic media. They help to identify whether a given solution is acidic or basic.

🔹 Types of Indicators:

Indicator	Acidic Solution	Basic Solution
Litmus	Blue → Red	Red → Blue
Phenolphthalein	Colorless	Pink
Methyl Orange	Red	Yellow

✅ Natural Indicators: Litmus (from lichens), turmeric, red cabbage

✅ Synthetic Indicators: Phenolphthalein, Methyl orange

🔁 5. Acid-Base Reactions with Indicators (Visual Table)

Substance	Effect on Litmus	Effect on Phenolphthalein	Effect on Methyl Orange
Hydrochloric acid	Blue → Red	Colorless	Red
Sodium hydroxide	Red → Blue	Pink	Yellow
Lemon juice	Blue → Red	Colorless	Red
Soap solution	Red → Blue	Pink	Yellow

🎯 6. Key Scientific Concepts

• Arrhenius Theory:

  • Acids: Produce H⁺ in water.

  • Bases: Produce OH⁻ in water.

• Neutralisation Reaction:

  • Acid + Base → Salt + Water
    HCl + NaOH → NaCl + H₂O

📚 7. Misconceptions Cleared

• Acids are not always dangerous. Many are edible and found in food (like citric acid).

• Not all substances that feel slippery are bases (e.g., oil is slippery but not basic).

• Indicators are tools, not proof of concentration or strength—just the presence of H⁺ or OH⁻ ions.

🧠 8. Thinking Corner

Why do ant bites cause a burning sensation, and how does applying baking soda help?

Answer: Ant sting injects formic acid (acidic). Baking soda (a base) neutralizes the acid, relieving the pain.

🧪 Hands-on Activity (Home/Lab)

Name: Testing Acids and Bases Using Indicators
Materials: Red/blue litmus paper, lemon juice, baking soda solution, soap, vinegar, shampoo
Procedure: Dip litmus in each solution and record color changes
Conclusion: Identify acidic and basic substances

📌 Summary Points

• Acids and bases are classified based on the release of H⁺ and OH⁻ ions respectively.

• Physical characteristics help in identification but are not to be tested manually.

• Indicators help detect the acidic or basic nature of substances.

• Some substances can behave as acids or bases depending on the environment (amphoteric substances – to be studied later).

Expressing Concentration of Solutions

 

📘 Expressing Concentration of Solutions

The composition of a solution refers to the relative amount of solute and solvent present in it. This composition can be expressed in two main ways:

• Qualitatively

• Quantitatively

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🔹 Qualitative Description

This is a non-numerical way of expressing the concentration. For example:

• A dilute solution has a relatively small amount of solute.

• A concentrated solution has a relatively large amount of solute.

However, qualitative terms are vague and can lead to confusion, especially in scientific or industrial applications. Hence, quantitative expressions of concentration are preferred.

🔹 Quantitative Description

Several methods are used to express the concentration of a solution quantitatively:

(i) Mass Percentage (% w/w)

Definition:

$$\text{Mass \% of a component} = \left( \frac{\text{Mass of the component in the solution}}{\text{Total mass of the solution}} \right) \times 100$$

Example:
A solution of 10% glucose by mass means:

• 10 g of glucose (solute)

• 90 g of water (solvent)
  Total mass = 100 g

Application:
Mass percentage is commonly used in industrial applications.
E.g., Commercial bleaching solution contains 3.62% NaOCl (sodium hypochlorite) by mass.

(ii) Volume Percentage (% v/v)

Definition:

$$\text{Volume \% of a component} = \left( \frac{\text{Volume of the component}}{\text{Total volume of the solution}} \right) \times 100$$

Example:
A 10% ethanol solution by volume means:

• 10 mL ethanol

• Water added to make the total volume = 100 mL

Application:
Used for liquid–liquid solutions like:

• Antifreeze: A 35% (v/v) ethylene glycol solution lowers the freezing point of water to 255.4 K (-17.6°C).

(iii) Mass by Volume Percentage (% w/v)

Definition:

$$\text{Mass by Volume \%} = \left( \frac{\text{Mass of solute in grams}}{\text{Volume of solution in mL}} \right) \times 100$$

Application:
Commonly used in medicine and pharmacy.
E.g., A 5% w/v solution means 5 g of solute in 100 mL of solution.

(iv) Parts per Million (ppm)

Definition:

$$\text{ppm} = \left( \frac{\text{Number of parts of the component}}{\text{Total number of parts of all components}} \right) \times 10^6$$

Application:
Used for very low concentrations (trace quantities), especially in:

• Environmental chemistry

• Pollution analysis
  E.g., Sea water (1030 g) contains 6 × 10⁻³ g of dissolved O₂

$$\text{ppm} = \frac{6 \times 10^{-3}}{1030} \times 10^6 \approx 5.8\ ppm$$

Variants: ppm can be:

• mass to mass

• volume to volume

• mass to volume

(v) Mole Fraction (χ)

Definition:

$$\chi_i = \frac{n_i}{\sum n_i}$$

Where:

• $n_i$ = number of moles of component i

• $\sum n_i$ = total moles of all components

In a binary mixture of components A and B:

$$\chi_A = \frac{n_A}{n_A + n_B}, \quad \chi_B = \frac{n_B}{n_A + n_B}$$ $$\chi_A + \chi_B = 1$$

Application:
Useful in calculating vapor pressure, colligative properties, and gas mixture compositions.

(vi) Molarity (M)

Definition:

$$\text{Molarity (M)} = \frac{\text{Moles of solute}}{\text{Volume of solution in litre}}$$

Example:
0.25 M NaOH means 0.25 mol of NaOH is dissolved in 1 litre of solution.

Note:
Temperature-dependent because volume changes with temperature.

(vii) Molality (m)

Definition:

$$\text{Molality (m)} = \frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$$

Example:
A 1.00 m solution of KCl means 1 mol (74.5 g) of KCl is dissolved in 1 kg of water.

Note:
Independent of temperature because it involves mass, not volume.

📝 Summary Table

Expression	Symbol	Formula	Depends on Temperature?	Typical Use
Mass % (w/w)	—	$\frac{\text{Mass of solute}}{\text{Total mass of solution}} \times 100$	❌ No	Industrial mixtures
Volume % (v/v)	—	$\frac{\text{Volume of solute}}{\text{Total volume of solution}} \times 100$	✅ Yes	Liquid-liquid solutions
Mass/Volume % (w/v)	—	$\frac{\text{Mass of solute}}{\text{Volume of solution in mL}} \times 100$	✅ Yes	Medical & pharmaceutical use
Parts per million (ppm)		$\frac{\text{Part of solute}}{\text{Total parts of solution}} \times 10^6$	❌ No	Environmental concentrations
Mole fraction	$\chi$	$\frac{\text{Moles of component}}{\text{Total moles of solution}}$	❌ No	Gas mixtures, vapor pressure
Molarity	M	$\frac{\text{Moles of solute}}{\text{Volume of solution in litres}}$	✅ Yes	Titrations, lab solutions
Molality	m	$\frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$	❌ No	Colligative property calculations

🔍 Conclusion

• Each unit has its specific application depending on the accuracy, temperature dependence, and nature of the solution.

• Mass %, ppm, mole fraction, and molality are independent of temperature, making them more reliable in changing environments.

• Molarity, while commonly used, must be handled carefully with temperature-sensitive systems.

Azeotropes – Explained for Class 12 (CBSE Chemistry)

 

📘 Azeotropes – Explained for Class 12 (CBSE Chemistry)

🔹 What is an Azeotrope?

An azeotrope is a mixture of two (or more) liquids that boils at a constant temperature and behaves like a pure substance during boiling.

🧪 Definition:

An azeotrope is a binary mixture that boils at a constant temperature and produces vapour with the same composition as the liquid.

This means you cannot separate the two components by simple distillation — they behave as one compound at that specific composition and boiling point.

🔍 Why does this happen?

Normally, during boiling, the composition of vapour differs from that of liquid (more volatile component escapes faster). But in an azeotropic mixture, this difference disappears at a certain fixed composition — the vapour formed has the same ratio of both components as in the liquid. Hence, no further separation by boiling is possible.

🌡️ Types of Azeotropes

Azeotropes are mainly of two types based on their boiling points in comparison to their pure components:

(i) Minimum Boiling Azeotropes

• These mixtures boil at a lower temperature than either of the pure components.

• They show positive deviation from Raoult’s Law (the components escape more easily due to weaker interactions).

• The total vapour pressure is higher, so they boil at a lower temperature.

🧪 Example:

• Ethanol (95%) + Water (5%)

• Boiling Point ≈ 351 K (78°C)

• This mixture forms an azeotrope that cannot be separated further by simple distillation.

🧠 Reason:

Ethanol and water form weaker hydrogen bonds in the mixture than in the pure liquids, so they escape more easily → higher vapour pressure → lower boiling point.

(ii) Maximum Boiling Azeotropes

• These mixtures boil at a higher temperature than either of the pure components.

• They show negative deviation from Raoult’s Law (the components strongly attract each other).

• The total vapour pressure is lower, so they boil at a higher temperature.

🧪 Example:

• Nitric acid (68%) + Water (32%)

• Boiling Point ≈ 393.5 K (120.5°C)

🧠 Reason:

Strong hydrogen bonding between nitric acid and water lowers the escaping tendency → lower vapour pressure → higher boiling point.

📊 Comparison Table:

Type	Boiling Point	Deviation from Raoult's Law	Vapour Pressure	Example
Minimum Boiling Azeotrope	Lower than components	Positive Deviation	High	Ethanol + Water (95%)
Maximum Boiling Azeotrope	Higher than components	Negative Deviation	Low	Nitric Acid + Water (68%)

🎓 Real-life Importance

• Azeotropes limit the extent of separation by distillation.

• Special methods like azeotropic distillation or adding third components (entrainers) are used in industries to break azeotropes.

• Ethanol production: 95% ethanol + 5% water azeotrope is common in alcohol distillation.

💡 Analogy to Understand

Think of azeotropes as inseparable couples:

• In a normal couple (mixture), one partner might leave early (more volatile component distills first).

• In an azeotrope, both are so balanced in their interactions that they leave together, hand in hand, at the same time (same composition in vapour and liquid).

Raoult’s Law and Henry’s Law – A Comparative Study

 

1. Raoult’s Law

Raoult’s law is applicable to solutions containing volatile components. It states:

"The partial vapour pressure of each volatile component in a solution is directly proportional to its mole fraction."

Mathematical Expression:

For a component 1 in a binary solution:

$$p_1 = x_1 \cdot p_1^0$$

Where:

• $p_1$ = Partial vapour pressure of component 1 in the solution

• $x_1$ = Mole fraction of component 1 in the solution

• $p_1^0$ = Vapour pressure of pure component 1

2. Henry’s Law

Henry’s Law is applicable to gases dissolved in liquids. It states:

"The partial pressure of the gas (p) in the vapour phase is directly proportional to its mole fraction (x) in the solution."

Mathematical Expression:

$$p = K_H \cdot x$$

Where:

• $p$ = Partial pressure of the gas

• $x$ = Mole fraction of the gas in solution

• $K_H$ = Henry’s Law constant

3. Raoult’s Law as a Special Case of Henry’s Law

• Both laws show a direct proportionality between partial pressure and mole fraction.

• In Raoult’s law, the constant of proportionality is $p_1^0$, while in Henry’s law it is $K_H$.

• Hence, Raoult’s law can be considered a special case of Henry’s law when $K_H = p_1^0$.

4. Vapour Pressure of Solutions

(a) Pure Solvent

• Molecules escape from the surface and exert vapour pressure at equilibrium.

(b) Solution with a Non-Volatile Solute

• When a non-volatile solute is added, the vapour pressure decreases.

• Fewer solvent molecules are available at the surface to escape into vapour phase.

Key Point:

• The reduction in vapour pressure depends only on the quantity of solute and not on its nature.

5. Ideal and Non-Ideal Solutions

(A) Ideal Solutions

Definition: Solutions that obey Raoult’s law over the entire concentration range.

Characteristics:

• $\Delta_{\text{mix}} H = 0$ (No heat is evolved or absorbed)

• $\Delta_{\text{mix}} V = 0$ (No change in volume on mixing)

• Intermolecular forces: A–A ≈ B–B ≈ A–B

Examples:

• Benzene and Toluene

• n-Hexane and n-Heptane

(B) Non-Ideal Solutions

Definition: Solutions that deviate from Raoult’s law.

They show either:

(i) Positive Deviation

• $p_{\text{solution}} > p_{\text{expected}}$

• A–B interactions < A–A and B–B interactions

• More molecules escape → Higher vapour pressure

Examples:

• Ethanol + Acetone

• Acetone + Carbon Disulphide

(ii) Negative Deviation

• $p_{\text{solution}} < p_{\text{expected}}$

• A–B interactions > A–A and B–B interactions

• Fewer molecules escape → Lower vapour pressure

Examples:

• Chloroform + Acetone (Hydrogen bonding between components)

• Phenol + Aniline

6. Azeotropes

Definition: Binary mixtures that boil at a constant temperature and whose liquid and vapour phases have the same composition.

Types of Azeotropes:

(i) Minimum Boiling Azeotropes

• Show positive deviation

• Boil at a lower temperature than both components

Example:

• Ethanol (95%) + Water → Boiling point: ~351 K

(ii) Maximum Boiling Azeotropes

• Show negative deviation

• Boil at a higher temperature than both components

Example:

• Nitric acid (68%) + Water → Boiling point: 393.5 K

7. Graphical Representations

(i) Raoult’s Law for Ideal Solutions

• Linear plot of vapour pressure vs. mole fraction

(ii) Non-Ideal Solutions

• Positive Deviation: Upward curve (convex)

• Negative Deviation: Downward curve (concave)

Summary Table

Type	Deviation	A-B Interaction	Vapour Pressure	Example
Ideal Solution	None	A–B ≈ A–A ≈ B–B	Matches Raoult’s law	Benzene + Toluene
Non-Ideal (+ve)	Positive	A–B < A–A or B–B	Greater than expected	Ethanol + Acetone
Non-Ideal (–ve)	Negative	A–B > A–A or B–B	Less than expected	Chloroform + Acetone
Min. Boiling Azeotrope	Positive	Weak A–B	Boils below both liquids	Ethanol (95%) + Water
Max. Boiling Azeotrope	Negative	Strong A–B	Boils above both liquids	Nitric Acid (68%) + Water

Conclusion

Raoult’s law and Henry’s law both describe the relationship between vapour pressure and mole fraction, although applicable in different contexts (liquids vs. gases). Understanding deviations from Raoult’s law helps us classify solutions as ideal or non-ideal and explains the formation of azeotropes, which are critical in processes like distillation.