ANAND CLASSES Study Material and Notes to explore all Laws of Chemical Combination (Law of Conservation of Mass, Law of Definite (Constant) Proportion, Law of Multiple Proportions, Law of Reciprocal (Equivalent) Proportions, Gay-Lussac’s Law of Gaseous Volumes) with clear definitions, solved examples, and a free PDF download. Essential for Class 11, JEE, NEET, and CBSE exam preparation.
Law of Conservation of Mass
Proposed by: Antoine Lavoisier
Statement:
“Matter is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.”
🧪 Example:
Reaction between AgNO3 and KI:
$$\text{AgNO}_3 (aq) + \text{KI} (aq) \rightarrow \text{AgI} (s) + \text{KNO}_3 (aq) $$
$$\text{Mass of AgNO}_3 + \text{Mass of KI} = \text{Mass of AgI} + \text{Mass of KNO}_3$$
Law of Definite (Constant) Proportion
Proposed by: Joseph Proust
Statement:
“A chemical compound always contains the same elements in a fixed ratio by mass, regardless of how it was prepared.”
🧪 Example:
Carbon dioxide (CO2) formation from different sources:
- From calcium carbonate:
$$\text{CaCO}_3 \xrightarrow{\Delta} \text{CaO} + \text{CO}_2$$
- From sodium bicarbonate:
$$2\text{NaHCO}_3 \xrightarrow{\Delta} \text{Na}_2\text{CO}_3 + \text{H}_2\text{O} + \text{CO}_2$$
In both cases: $$\text{Ratio of mass of C:O} = 12:32 = 3:8$$
Law of Multiple Proportions
Proposed by: John Dalton
Verified by: Berzelius
Statement:
“When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of simple whole numbers.”
🧪 Example:
Nitrogen forms the following oxides:
$$\text{N}_2\text{O}, \text{NO}, \text{N}_2\text{O}_3, \text{N}_2\text{O}_4, \text{N}_2\text{O}_5$$
With 28g nitrogen, the oxygen mass is: $$16:32:48:64:80 = 1:2:3:4:5$$
🔷 Law of Reciprocal (Equivalent) Proportions
📜 Proposed By: Richter (1792)
He was among the first to recognize that chemical reactions occur in definite and fixed proportions by mass, which laid the foundation for stoichiometry.
🧠 Statement of the Law:
“The masses of two elements which separately combine with a fixed mass of a third element are either the same or in a simple whole-number ratio when they combine with each other.”
🧪 Detailed Example:
Let’s take Hydrogen (H), Sulphur (S), and Oxygen (O) to illustrate this law.
🔹 Step 1: How H combines with S
$$\text{2g H} + \text{32g S} \longrightarrow \text{H}_2\text{S}$$
➡️ So, 2g of hydrogen combines with 32g of sulphur.
🔹 Step 2: How H combines with O
$$\text{2g H} + \text{16g O} \longrightarrow \text{H}_2\text{O} $$
➡️ So, 2g of hydrogen also combines with 16g of oxygen.
🔹 Step 3: Now combine S and O directly
$$\text{32g S} + \text{32g O} \longrightarrow \text{SO}_2 $$
➡️ Here, 32g of sulphur combines with 32g of oxygen.
📊 Now apply the Law:
Let’s compare how S and O combine with H and then with each other:
- S combines with 2g H → 32g S
- O combines with 2g H → 16g O
So, S and O individually combine with equal mass (2g) of H.
Now check the ratio in which S and O combine with each other: $$\frac{32}{16} : \frac{32}{32} = 2 : 1$$
✔️ A simple whole number ratio → Law is verified!
📚 Concept Behind the Law:
Think of it like this:
If A combines with B and A combines with C, then B and C should also combine in the same or simple multiple ratio when forming a new compound.
This law indirectly supports the idea that atoms combine in whole numbers (as later supported by Dalton’s Atomic Theory and Avogadro’s Law).
💡 Real-Life Example
- Carbon (C) combines with Oxygen (O) to form both CO and CO₂.
- Carbon also combines with Hydrogen (H) to form CH₄.
- Now compare how H and O combine with C, and then see how they combine together in H₂O → You’ll observe this law again!
Gay-Lussac’s Law of Gaseous Volumes
Proposed by: Gay-Lussac
Statement:
“When gases react, they do so in simple whole number ratios by volume under similar temperature and pressure.”
🧪 Example 1:
$$\text{H}_2 + \frac{1}{2} \text{O}_2 \rightarrow \text{H}_2\text{O}$$
$$1 \text{ volume H}_2 + 0.5 \text{ volume O}_2 \rightarrow 1 \text{ volume H}_2\text{O}$$
🧪 Example 2:
$$\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3$$
Volumes: $1 : 3 : 2$
🌟 Avogadro’s Law – Explained
🔬 What is Avogadro’s Law?
In 1811, Amedeo Avogadro, an Italian scientist, proposed a fundamental idea:
“Equal volumes of all gases, at the same temperature and pressure, contain an equal number of molecules.”
📘 Statement:
If:
- Volume V is directly proportional to the number of moles n,
- At constant Temperature T and Pressure P,
Then :
V ∝ n(at constant T and P)
or $$\frac{V_1}{n_1} = \frac{V_2}{n_2}$$
🧠 Understanding with Example: Formation of Water
Consider the reaction: $$2 \text{H}_2 (g) + \text{O}_2 (g) \rightarrow 2 \text{H}_2\text{O} (g)$$
🔹 Volume Relationship:
- 2 volumes of Hydrogen
- 1 volume of Oxygen
- Form 2 volumes of Water vapor
➡️ This volume ratio suggests a fixed number of gas particles (molecules) per volume if T and P are constant.
🧪 Why It Was Revolutionary
At the time, scientists believed:
- Atoms were indivisible
- Same atoms couldn’t combine (like H + H)
- Molecules like $\text{H}_2,$ $\text{O}_2$ didn’t exist
Avogadro introduced the distinction between atoms and molecules, suggesting that:
- Gases like hydrogen and oxygen are diatomic molecules: $$\text{H}_2, \text{O}_2$$
This helped explain:
- Why 2 volumes of $H_2$ react with 1 volume of $O_2$ to give 2 volumes of $H_2O$.
✨ Key Concepts
- Molar Volume:
At STP (0°C, 1 atm), 1 mole of any gas occupies 22.4 L - Avogadro’s Number: $N_A = 6.022 \times 10^{23} \text{ molecules/mole}$
- Volume and Mole Relationship: $$\frac{22.4 \text{ L}}{1 \text{ mole}} = \frac{V}{n} \Rightarrow V = n \times 22.4 \text{ L}$$
🧠 Conceptual Q&A
Q1. Why does 1 mole of oxygen occupy the same volume as 1 mole of hydrogen at STP?
A: Because Avogadro’s law states that equal moles of gases occupy equal volumes under the same conditions of temperature and pressure.
Q2. What is the significance of Avogadro’s Law in stoichiometry?
A: It helps in converting volumes of gases to moles in chemical reactions, making calculations easier.
💡 Do You Know?
- Water always contains hydrogen and oxygen in a mass ratio of 1:8.
- The law of conservation of mass fails in nuclear reactions.
- Gay-Lussac’s Law laid the path to Avogadro’s Hypothesis.
- Mixtures do not obey the Law of Definite Proportion.
❓ FAQs
Q1: Why doesn’t the Law of Conservation of Mass apply to nuclear reactions?
Ans: Because mass converts to energy, violating mass conservation.
Q2: Can Gay-Lussac’s Law be applied to liquids?
Ans: No, it applies only to gases because it is volume-based.
Q3: Which law supports Dalton’s atomic theory?
Ans: Law of Multiple Proportions
🧠 Conceptual Questions
Q1: Which law explains the fixed ratio of hydrogen and oxygen in water?
Ans: Law of Definite Proportion
Q2: If 14g N combines with 16g and 48g O, which law is this?
Ans: Law of Multiple Proportion
🎯 MCQs with Explanation
Q1: A chemical compound always has the same elements in the same mass ratio. This is:
A. Conservation of Mass
B. Definite Proportion ✅
C. Reciprocal Proportion
D. Gaseous Volumes
✔️ Explanation: Law of Definite Proportion deals with fixed ratios by mass.
Q2: Who proposed the Law of Conservation of Mass?
A. Dalton
B. Proust
C. Lavoisier ✅
D. Berzelius
✔️ Explanation: Lavoisier is known as the “Father of Modern Chemistry”.
📘 Worksheet
Q1: 10g of CaCO3 decomposes and gives 5.6g of CaO. Find mass of CO2. Ans: 10−5.6=4.4 g
Q2: A gas obeys simple volume ratios in reactions. Which law is this?
Ans: Gay-Lussac’s Law
🔄 Quick Revision Points
- 🔵 Law of Conservation of Mass: Mass remains constant.
- 🟡 Law of Definite Proportion: Fixed composition by mass.
- 🔴 Law of Multiple Proportion: Whole number ratios of mass.
- 🟢 Law of Reciprocal Proportion: Cross-combination ratios.
- 🟣 Gay-Lussac’s Law: Gases react in volume ratios.
