🧪 Unit 1: Some Basic Concepts of Chemistry

 

🔶 1.1 Importance of Chemistry

 

🔹 Central Role of Chemistry:

• Chemistry plays a central role in science and is interlinked with physics, biology, geology, etc.
• It helps explain natural phenomena, like weather patterns, brain function, and even computer operations.

 

🔹 Role in Daily Life:

• Used in manufacturing of:
  🧫 Fertilisers, 🧪 Acids, 🧼 Soaps, 🧬 Drugs, 🧲 Polymers, 🧴 Cosmetics, and more.

 

🔹 Economic and Social Importance:

• Drives national economic growth.
• Improves quality of life through:
  • Food production
  • Healthcare products
  • New technologies (e.g., superconductors, optical fibres)

 

🔹 Environmental Applications:

• Safer alternatives to ozone-depleting CFCs created.
• Chemistry contributes to understanding and managing greenhouse gases like CO₂ and CH₄.

 

🔹 Future Needs:

• India needs creative and talented chemists to tackle:
  • Biochemical reactions
  • Enzyme-based production
  • New synthetic materials

 

🔶 1.2 Nature of Matter

 

🔹 What is Matter?

• Anything that has mass and occupies space.
• Examples: Book, air, water, living organisms.

 

🧊 1.2.1 States of Matter

State	Volume	Shape	Particle Arrangement
Solid	Definite	Definite	Tightly packed, ordered
Liquid	Definite	Indefinite	Close but free to move
Gas	Indefinite	Indefinite	Far apart, move freely

🔄 Interconversion between states:
Solid ⇌ Liquid ⇌ Gas (via change in temperature/pressure)

 

🧪 1.2.2 Classification of Matter

 

🔹 Based on Composition:

1. Pure Substances
   • Fixed composition
   • Elements or Compounds
   • Cannot be separated by physical methods
   • Examples: Glucose, Water, Copper
2. Mixtures
   • Variable composition
   • Can be homogeneous (uniform) or heterogeneous (non-uniform)
   • Separable by physical methods (e.g., filtration, distillation)

Type	Example	Key Feature
Homogeneous	Sugar solution, air	Uniform throughout
Heterogeneous	Soil, salad, milk	Components visible/separable

 

⚛️ Elements vs Compounds:

Aspect	Element	Compound
Nature	One type of atom	Two or more types of atoms
Simplification	Cannot be broken further	Can be broken into simpler elements
Example	Oxygen (O₂), Iron (Fe)	Water (H₂O), Carbon Dioxide (CO₂)

Perfect! Let’s continue with the next section:

 

🔷 1.3 Properties of Matter and Their Measurement

 

🧪 Physical Properties

These are measurable without changing the identity of the substance.

🧾 Examples:

• Colour
• Odour
• Melting point
• Boiling point
• Density
• State of matter

 

⚗️ Chemical Properties

These describe how a substance reacts with other substances, i.e., changes into something new.

🧾 Examples:

• Reactivity with acid
• Combustibility
• Rusting of iron
• Oxidation of copper

 

📐 1.3.1 Measurement

🧠 Measurement = Quantitative observation
It involves two parts:

• A number
• A unit (standard of comparison)

🧾 Example:

• Length = 2 metres
• Mass = 1.5 kilograms

 

📏 SI Units (Système Internationale d’Unités)

Quantity	SI Unit	Symbol
Length	metre	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Electric current	ampere	A
Luminous intensity	candela	cd

 

📊 1.3.2 Base Units and Derived Units

• Base Units → 7 fundamental units (like metre, kilogram, second, etc.)
• Derived Units → Formed by combining base units

🧾 Example:

• Speed = distance/time = m/s
• Volume = m × m × m = m³
• Density = mass/volume = kg/m³

 

🧪 1.3.3 Mass and Weight

Concept	Description
Mass	Quantity of matter in a body
Weight	Force due to gravity on that mass (W = mg)
Mass	Remains constant
Weight	Varies with gravitational field

🧠 Mass is measured in kilograms, weight in newtons

 

🌡️ 1.3.4 Volume

• Volume = amount of space occupied
• SI unit = cubic metre (m³)
• Often expressed in:
  • cm³ = 1 mL
  • dm³ = 1 L

🧾 1 L = 1000 mL = 1000 cm³
🧾 1 m³ = 1000 L

 

🌡️ 1.3.5 Temperature

Scale	Symbol	Freezing Point of Water	Boiling Point of Water
Celsius	°C	0°C	100°C
Fahrenheit	°F	32°F	212°F
Kelvin	K	273.15 K	373.15 K

📌 Kelvin is SI unit
🔁 Conversion formulas:

Awesome! Let’s now move forward with:

 

🟩 1.3.6 Density

 

📌 What is Density?

Density is the mass per unit volume of a substance.

 

📏 SI Unit:

• Kilogram per cubic metre (kg/m³)
• Commonly used: g/cm³ or g/mL

📌 Conversions:

• 1 g/cm³ = 1000 kg/m³
• 1 mL = 1 cm³

 

🧪 Example:

If an object has mass = 250 g and volume = 200 mL

 

🟦 1.4 The International System of Units

 

📌 Need for Standard Units:

To avoid confusion, scientists use a standard system of measurement → SI system

✅ It is universal, rational, and metric-based

 

🔹 7 SI Base Units (Revisited Table):

Base Quantity	SI Unit	Symbol
Length	metre	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Electric Current	ampere	A
Amount of Substance	mole	mol
Luminous Intensity	candela	cd

 

🔹 Some Derived Units:

Quantity	Formula	SI Unit
Area	length × breadth	m²
Volume	length × breadth × height	m³
Speed	distance / time	m/s
Density	mass / volume	kg/m³
Pressure	force / area	pascal (Pa) = N/m²
Energy	force × distance	joule (J)
Power	energy / time	watt (W)

Perfect! Let’s now move to:

 

🟨 1.5 Uncertainty in Measurement

 

📌 Why Uncertainty Arises:

• Every measurement contains some error or uncertainty.
• No instrument is perfectly precise, and personal error may occur while reading.

 

🔹 Accuracy vs. Precision

Term	Meaning
Accuracy	How close a measured value is to the true value
Precision	How close repeated measurements are to each other

🧠 A measurement can be precise but not accurate, or accurate but not precise.

 

🔍 Example:

• True value of a mass = 5.0 g
• If measured as: 5.3 g, 5.2 g, 5.3 g → Precise but not accurate
• If measured as: 5.0 g, 5.0 g, 5.0 g → Both accurate and precise ✅

 

⚠️ Types of Errors:

1. Systematic Errors
   • Occur consistently due to instrument or method issues.
   • Examples: Zero error in a balance, reaction time error.
2. Random Errors
   • Occur unpredictably, due to small uncontrollable variations.
   • Examples: Fluctuations in temperature, estimation in timing.
3. Gross Errors
   • Due to carelessness or mistake (e.g., reading scale wrong, calculation mistake)

 

🟩 1.6 Significant Figures

 

📌 What are Significant Figures?

They indicate the certainty and precision in a measured quantity.

👉 Includes:

• All certain digits + first uncertain digit

 

🧠 Rules for Counting Significant Figures:

Rule #	Description	Example
1	All non-zero digits are significant	237 → 3 sig. fig.
2	Zeros between non-zero digits are significant	1007 → 4
3	Leading zeros (before first non-zero) are not	0.0045 → 2
4	Trailing zeros in a number with decimal are significant	2.500 → 4
5	Trailing zeros in a number without decimal are not significant	2500 → 2
6	Exact numbers (like 2 in “2πr”) have infinite significant figures	–

 

🧪 Examples:

Number	Significant Figures
0.070	2
6.002	4
400	1
400.0	4

Great! Now let’s cover:

 

🟦 1.6.1 Rules for Arithmetic Operations with Significant Figures

 

➕ (A) Addition and Subtraction

📌 Rule:
The result should have the same number of decimal places as the number with the least decimal places.

 

🧪 Example:

• 18.0 has 1 decimal place, so
  ✅ Final Answer = 31.1

 

✖️ (B) Multiplication and Division

📌 Rule:
The result should have the same number of significant figures as the number with the least significant figures.

 

🧪 Example:

• 4.56 → 3 significant figures
• 1.4 → 2 significant figures ⇒ So, final answer should have 2 significant figures

 

🟨 1.6.2 Rounding Off the Numbers

 

📌 Rules for Rounding:

1. If the digit to be dropped is < 5 → leave the preceding digit unchanged.
   ✅ Example: 3.74 → 3.7
2. If the digit to be dropped is > 5 → increase the preceding digit by 1.
   ✅ Example: 3.78 → 3.8
3. If the digit to be dropped is 5:
   • If the previous digit is even, leave it.
   • If it is odd, increase by 1.

📌 Examples:

• 3.75 → 3.8 (5 after odd digit → round up)
• 3.65 → 3.6 (5 after even digit → leave it)

Awesome! Let’s now continue with:

 

🟩 1.6.3 Determining Uncertainty in Calculated Results

 

📌 What is Uncertainty?

Every measured value has some uncertainty (error), especially when the value is used in calculations.

 

🧪 Example 1:

Length = cm
Breadth = cm

➡️ Area = cm²

% Uncertainty in length =
% Uncertainty in breadth =

➡️ Total % Uncertainty =
🧾 Final result = cm² ✅

 

📌 Tips for Multi-step Calculations:

🔹 Keep 1 extra digit in intermediate steps.
🔹 Round only in the final step to avoid cumulative rounding errors.

 

🟨 1.7 Laws of Chemical Combinations

These laws describe how elements combine to form compounds.

 

⚖️ 1.7.1 Law of Conservation of Mass

➡️ Mass is neither created nor destroyed during a chemical reaction.

📖 Proposed by Antoine Lavoisier (1789)

🧪 Example:

✔️ Total mass of products = Mass of reactants

 

⚖️ 1.7.2 Law of Definite Proportions (Constant Composition)

➡️ A given compound always contains the same elements in the same fixed ratio by mass, regardless of source.

📖 Proposed by Joseph Proust (1799)

🧪 Example:
Water (H₂O) always has H:O in 1:8 mass ratio
(2 g H + 16 g O = 18 g H₂O)

 

⚖️ 1.7.3 Law of Multiple Proportions

➡️ When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole number ratios.

📖 Proposed by John Dalton (1803)

🧪 Example:

• CO → 12 g C + 16 g O
• CO₂ → 12 g C + 32 g O
  Ratio of oxygen masses = 16:32 = 1:2

 

⚖️ 1.7.4 Gay Lussac’s Law of Gaseous Volumes

➡️ Gases react in simple volume ratios (at same T and P).

🧪 Example:
H₂ + ½ O₂ → H₂O
→ 2 volumes H₂ + 1 volume O₂ → 2 volumes H₂O vapor
Ratio: 2 : 1 : 2

 

⚖️ 1.7.5 Avogadro’s Law

➡️ Equal volumes of gases (at same T and P) contain equal number of molecules.

📌 Important for:

• Determining molar volume of gases
• Calculating atomic/molecular weights

Awesome! Let’s now proceed with:

 

🧪 1.8 Dalton’s Atomic Theory

 

📌 Proposed by:

John Dalton in 1808

 

🔷 Main Postulates:

1. Matter consists of indivisible atoms.
2. All atoms of a given element are identical in mass and properties.
3. Atoms of different elements are different in mass and properties.
4. Compounds are formed when atoms of different elements combine in fixed whole number ratios.
5. Chemical reactions involve rearrangement of atoms; atoms are neither created nor destroyed.

 

📌 Significance:

Dalton’s theory laid the foundation for modern chemistry, especially the laws of chemical combinations.

 

⚠️ Limitations:

• Atoms are divisible (protons, neutrons, electrons)
• Atoms of an element can have different masses (isotopes)
• Does not explain the structure or bonding of atoms

 

🧪 1.9 Atomic and Molecular Masses

 

🔹 1.9.1 Atomic Mass

Atomic mass of an element = Relative mass of its atom as compared to 1/12 of a carbon-12 atom.

📌 Modern unit: 1 amu = 1 unified mass unit (u)

 

🧠 Examples:

Element	Atomic Mass (u)
Hydrogen (H)	1.008
Oxygen (O)	16.00
Carbon (C)	12.00

 

🔹 1.9.2 Molecular Mass

Molecular mass = Sum of atomic masses of all atoms in a molecule.

 

🧪 Examples:

• H₂O = 2 × 1.008 (H) + 16.00 (O) = 18.016 u
• CO₂ = 12.00 (C) + 2 × 16.00 (O) = 44.00 u

Perfect! Let’s now continue with:

 

🧮 1.10 Mole Concept and Molar Masses

 

🔹 What is a Mole?

➡️ A mole is the amount of a substance that contains the same number of entities (atoms, molecules, ions, or particles) as there are in 12 grams of carbon-12.

📌 Avogadro’s Number (NA):

 

🔹 Key Relationships:

Quantity	Formula
Number of moles (n)
Number of particles
Volume of gas at STP

📌 STP: Standard Temperature and Pressure (0°C, 1 atm)

 

🧠 Example:

How many molecules are there in 18 g of water?

• Molar mass of H₂O = 18 g/mol
• Moles = 18 / 18 = 1 mol
• Molecules = = 6.022 × 10²³ molecules

 

🔹 Molar Mass:

The mass of one mole of a substance, equal to the molecular or atomic mass expressed in grams.

🧾 Examples:

• Molar mass of H₂O = 18 g/mol
• Molar mass of CO₂ = 44 g/mol

 

🧮 1.11 Percentage Composition

 

📌 What is it?

The percentage by mass of each element present in 100 g of a compound.

 

🔹 Formula:

 

🧪 Example: Water (H₂O)

• Molar mass = 18 g/mol
• Mass of H = 2 × 1 = 2 g
• Mass of O = 16 g

✅ Total = 100%

Awesome! Let’s complete the final sections of Chapter 1:

 

📏 1.12 Empirical and Molecular Formula

 

📌 Empirical Formula

The formula that shows the simplest whole number ratio of atoms in a compound.

 

📌 Molecular Formula

The formula that shows the actual number of atoms of each element in one molecule of the compound.

 

🔹 Relationship:

Where:

 

🧪 Example: Glucose (C₆H₁₂O₆)

• Empirical formula = CH₂O
• Empirical formula mass = 30
• Molar mass = 180

 

⚖️ 1.13 Stoichiometry and Stoichiometric Calculations

 

📌 What is Stoichiometry?

Stoichiometry is the quantitative relationship between reactants and products in a balanced chemical reaction.

 

🧪 Example Reaction:

Reactant/Product	Moles	Molecules
H₂	2
O₂	1
H₂O	2

 

🔹 Mass Relationships:

• 2 mol H₂ = 4 g
• 1 mol O₂ = 32 g
• 2 mol H₂O = 36 g

📌 Thus,
4 g H₂ + 32 g O₂ → 36 g H₂O ✅

 

🧪 Example Problem:

How many grams of CO₂ are produced when 5 g of carbon burns in excess oxygen?

• Reaction:
• Molar mass of C = 12 g/mol
• Moles of C = 5 / 12 ≈ 0.417 mol
• 1 mol C gives 1 mol CO₂ ⇒ 0.417 mol CO₂
• Molar mass of CO₂ = 44 g/mol

✅ Final Answer: 18.35 g CO₂