1. Importance of Chemistry
- Role of chemistry in daily life, industry, medicine, environment, etc.
🏠 1. Daily Life
- Food preservation: Use of preservatives like sodium benzoate.
- Cleaning agents: Soaps, detergents, and shampoos are chemical products.
- Cooking: Chemical reactions like Maillard reaction and baking soda action.
- Cosmetics: Lipsticks, creams, and perfumes are made using organic compounds.
🏭 2. Industry
- Petroleum industry: Produces fuels, plastics, and synthetic fibers.
- Fertilizer industry: Chemistry helps create urea, NPK fertilizers, etc.
- Cement and glass industries: Rely on chemical processes for manufacturing.
- Textile industry: Use of dyes, synthetic fibers like nylon and polyester.
💊 3. Medicine
- Drugs and vaccines: Painkillers (aspirin), antibiotics (penicillin), and vaccines are all chemistry-based.
- Diagnostics: Chemical indicators used in tests like blood glucose, pH.
- Pharmaceuticals: Development of new medicines and treatment methods.
🌍 4. Environment
- Pollution control: Chemistry helps develop scrubbers, filters, and biodegradable materials.
- Water purification: Use of chlorine, ozone, and filtration techniques.
- Climate studies: Chemistry is used in studying greenhouse gases and global warming.
- Recycling and waste management: Chemical methods are used to recycle plastics, metals, etc.
1. Law of Conservation of Mass
📘 Statement:
“Mass can neither be created nor destroyed in a chemical reaction.”
This means the total mass of the reactants = total mass of the products.
🧪 Example:
Let’s consider this reaction:
Hydrogen (H₂)+Oxygen (O₂)→Water (H₂O)
Suppose:
- 2 grams of hydrogen react with 16 grams of oxygen
- They form 18 grams of water
- So,
- Mass of reactants = 2 g + 16 g = 18 g
- Mass of products = 18 g
- → Hence, mass is conserved.
2. Law of Definite Proportions (or Constant Composition)
📘 Statement:
“A given compound always contains the same elements in the same fixed ratio by mass, regardless of its source or method of preparation.”
🧪 Example:
Let’s take water (H₂O) as an example.
- It always contains hydrogen and oxygen in the mass ratio of 1:8.
- That means:
- 2 grams of hydrogen combine with 16 grams of oxygen
- Ratio = 2 : 16 = 1 : 8
✅ So, whether you get water from:
- A river,
- A chemical lab, or
- Rain —
It always has hydrogen and oxygen in the 1:8 mass ratio.
3. Law of Multiple Proportions
📘 Statement:
“When two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.”
🧪 Example:
Let’s take carbon and oxygen.
They form two compounds:
- Carbon monoxide (CO)
- Carbon dioxide (CO₂)
Let’s fix mass of carbon = 12 g in both compounds:
- In CO: 12 g of carbon combines with 16 g of oxygen
- In CO₂: 12 g of carbon combines with 32 g of oxygen
Now, compare oxygen masses:
16 : 32 = 1 : 2 → a simple whole number ratio ✅
➡️ This proves the law of multiple proportions.
4. Gay-Lussac’s Law of Gaseous Volumes
📘 Statement:
“When gases react together, they do so in volumes which bear a simple whole number ratio to each other and to the volumes of the products (if gaseous), under the same temperature and pressure.”
🧪 Example:
Let’s take the reaction of hydrogen and oxygen forming water vapor: 2H2(g)+O2(g)→2H2O(g)
Let’s look at the volumes:
- 2 volumes of hydrogen
- 1 volume of oxygen
- produce 2 volumes of water vapor
So the ratio is:
2 : 1 : 2 — a simple whole number ratio.
This verifies Gay-Lussac’s Law.
5. Avogadro’s Law
📘 Statement:
“Equal volumes of all gases, at the same temperature and pressure, contain an equal number of molecules.”
🧪 Implication:
- 1 mole of any gas at STP (Standard Temperature and Pressure) occupies 22.4 L
- So, 1 L of H₂ gas has the same number of molecules as 1 L of O₂ gas under same conditions.
- 🧪 Example:
- Consider: H2(g)+Cl2(g)→2HCl(g)
- 1 volume of hydrogen
- 1 volume of chlorine
- gives 2 volumes of hydrogen chloride gas
- Since volumes are in a simple ratio, and Avogadro’s law tells us equal volumes → equal number of molecules, this reaction supports both Avogadro’s Law and Gay-Lussac’s Law.
- Dalton’s Atomic Theory (1808)
- 📘 Proposed by: John Dalton
- This theory was the first scientific attempt to explain the nature of matter using atoms.
Main Postulates of Dalton’s Atomic Theory:
- All matter is made up of tiny particles called atoms.
- Atoms are indivisible and indestructible (later proven partially wrong).
- All atoms of a given element are identical in mass and properties.
- Atoms of different elements have different masses and properties.
- Atoms combine in simple whole number ratios to form compounds.
- Atoms can neither be created nor destroyed in a chemical reaction.
- They are simply rearranged.
- The relative number and kinds of atoms are constant in a given compound.
Example:
Water (H₂O):
- 2 atoms of hydrogen + 1 atom of oxygen → combine in a fixed ratio (2:1)
- The type and number of atoms remain the same, as Dalton’s theory says.
- Limitations:
- Atoms are not indivisible (they are made of protons, neutrons, electrons).
- Isotopes and isobars contradict the idea that all atoms of an element are identical.
- Atomic Mass and Molecular Mass
1. Atomic Mass
📘 Definition:
The atomic mass of an element is the mass of a single atom, expressed in atomic mass units (amu or u).
1 amu is defined as 1/12th the mass of one carbon-12 atom.
Examples:
| Element | Atomic Mass (approx) |
|---|---|
| Hydrogen (H) | 1 u |
| Carbon (C) | 12 u |
| Oxygen (O) | 16 u |
| Sodium (Na) | 23 u |
2. Molecular Mass
📘 Definition:
The molecular mass of a compound is the sum of atomic masses of all atoms in a molecule.
🧪 Formula:
Molecular Mass=∑(Atomic masses of all atoms in the molecule)
Examples:
- Water (H₂O):
= (2 × 1) + (1 × 16) = 18 u - Carbon Dioxide (CO₂):
= (1 × 12) + (2 × 16) = 44 u - Methane (CH₄):
= (1 × 12) + (4 × 1) = 16 u
Mole Concept
📘 Definition:
A mole is the amount of a substance that contains as many particles (atoms, molecules, ions) as there are atoms in 12 grams of carbon-12.
✅ That number is called Avogadro’s number (NA)
NA=6.022×1023 particles/mol
💡 Key Relationships in Mole Concept:
| Quantity | Formula |
|---|---|
| Number of moles (n) | Molar mass (g/mol) ________________ Given mass (g) |
| Number of particles | Moles ×6.022×1023 |
| Mass from moles | Moles × Molar mass |
| Volume at STP | 1 mole of gas = 22.4 L at STP (0°C, 1 atm) |
🧪 Example
1.How many moles are there in 18 g of water (H₂O)?
Molar mass of H₂O = 18 g/mol
n=18/18=1 mole
2.How many molecules are there in 1 mole of CO₂?
6.022×1023 Molecules
3.Volume occupied by 2 moles of oxygen gas at STP?
= 2×22.4=44.8 L
Percentage Composition
📘 Definition:
Percentage composition of a compound tells us the percentage by mass of each element present in that compound.
🧮 Formula:
% of element=
(Mass of the element in 1 mole of compound /Molar mass of compound) X 100
🧪 Example:
Calculate the percentage composition of water (H₂O):
- Molar mass of H₂O = 18 g/mol
- Mass of hydrogen in 1 mole = 2 g
- Mass of oxygen in 1 mole = 16 g
%H = 2/18×100=11.11%
%O = 16/18×100 = 88.89%
🧮 1. Empirical Formula
📘 Definition:
The empirical formula shows the simplest whole-number ratio of atoms of each element in a compound.
Example:
- Hydrogen peroxide (H₂O₂)
- Empirical formula: HO
- Molecular formula: H₂O₂
- Glucose (C₆H₁₂O₆)
- Empirical formula: CH₂O
- Molecular formula: C₆H₁₂O₆
🧬 2. Molecular Formula
📘 Definition:
The molecular formula shows the actual number of atoms of each element in a molecule.
🧮 Formula to find Molecular Formula:
Molecular Formula=n X Empirical Formula
Where:
n = Molar Mass of compound / Empirical formula mass
🔍 Example:
If empirical formula = CH₂,
Empirical formula mass = 12 + (2×1) = 14 u
Given molar mass = 28 u
n = 28/14= 2
So, Molecular formula = (CH₂) × 2 = C₂H₄
*Stoichiometry and Stoichiometric Calculations
📘 Definition:
Stoichiometry is the study of the quantitative relationships (mass or volume) between reactants and products in a balanced chemical equation.
🧪 Example Reaction:
2H2+O2→2H2O
This tells us:
- 2 moles of H₂ react with 1 mole of O₂ to produce 2 moles of H₂O.
Or in mass:
2×2=4 g H₂
1×32=32 g O₂
→ gives 2×18=36 g
So the mass ratio is:
4 g+32 g=36 g
🔢 Typical Stoichiometric Problems Involve:
- Finding the mass of products/reactants
- Limiting reagent calculations
- Finding volume of gases at STP
- Number of moles or particles
Q: How many grams of water are formed when 4 grams of hydrogen reacts with excess oxygen?
Step 1: Balanced Equation
2H2+O2→2H2O
Step 2: Molar Mass
- H₂ = 2 g/mol
- H₂O = 18 g/mol
Step 3: Use ratio
From equation: 4 g H₂ gives 2 × 18 = 36 g H₂O
Answer: 36 g of water is formed
Limiting Reagent
📘 Definition:
The limiting reagent is the reactant that gets completely used up first in a chemical reaction, thus limiting the amount of product formed.
🧪 Why Important?
In many reactions, one reactant is in excess, and the other limits how much product is made.
Once the limiting reagent is used up the reaction stops, even if other reactants are still present.
🔍 Example:
Reaction: 2H2+O2→2H2O
Suppose you have:
- 5 moles of H₂
- 2 moles of O₂
- mole ratio required: 2 mol H₂ : 1 mol O₂
- So for 2 mol O₂, you’d need 4 mol H₂. You have 5 mol H₂, which is more than needed.
- Thus, O₂ is the limiting reagent, and it will decide the amount of H₂O formed.
🧮 How to Identify Limiting Reagent:
- Write the balanced equation
- Convert given masses to moles
- Divide available moles by the coefficient in the balanced equation
- Smaller result → Limiting reagent
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