Revision: Some Basic Concepts of Chemistry Chemistry Science (English Medium) Class 11 CBSE

The weight of an object is defined as the force with which the earth attracts the object.

Mass is the amount of matter present in the object. The SI unit of mass is kg.

The measured value of a physical quantity denoting the number of digits in which we have confidence — where a larger number indicates greater accuracy of measurement — is called significant figures.

An atom is the smallest particle of an element that can take part in a chemical reaction; however, it may or may not exist independently. 

A molecule is the smallest particle of an element or a compound that can exist by itself; it never breaks up except for taking part in a chemical reaction.

Relative atomic mass is defined as the ratio of the average atomic mass to the unified atomic mass unit.

Relative atomic mass (A_r) = `"Average mass of the atom"/"Unified atomic mass"`

The mass of a single atom of an element is called the atomic mass.

The weighted average of the masses of all its isotopes in a sample of that element is called the average molecular mass.

The sum of the atomic masses of all atoms in a molecule is called the molecular mass.

Avogadro’s number is defined as the number of atoms present in 12g of ₆C¹² isotope i.e. 6.023 × 10²³ atoms.

"The relative atomic mass or atomic weight of an element is the number of times one atom of the element is heavier than `1/12` times of the mass of an atom of carbon - 12".
Relative atomic mass = Mass of 1 atom of the element `1/12` of the mass of one C¹² atom.

Mole is the amount of a substance containing elementary particles like atoms, molecules or ions in 12 g of carbon - 12.

One mole of any gaseous molecules occupies 22.4 dm³ (litre) or 22400 cm³ (ml) at standard temperature and pressure (STP). This volume is known as the molar volume.

The relative molecular mass of a compound is the number that represents how many times one molecule of the substance is heavier than `1/12` of the mass of an atom of carbon ₆C¹².

The quantity of the element which weighs equal to its gram atomic mass is called one gram atom of that element.

A mole is the amount of pure substance containing the same number of chemical units as there are atoms in exactly 12 grams of carbon -12.

Avogadro's number is defined as the number of atoms present in 12 g (gram atomic mass) of C-12 isotope, i.e., 6·022 x 1023 atoms.

OR

Avogadro's number is the number of elementary units, i.e., atoms, ions or molecules present in one mole of a substance. It is denoted by N_A.

Percentage composition of a compound, is the percentage by weight of each element present in it.

The limiting reagent is the reactant that is completely consumed in a chemical reaction and determines the maximum amount of product that can be formed.

• Unsaturated solution: If the amount of solute contained in a solution is less than the saturation level, it is called an unsaturated solution. (till it is dissolving).

• Saturated solution: When no more solute can be dissolved in a solution at a given temperature, it is called a saturated solution.

• Solubility: The amount of the solute present in the saturated solution at this temperature is called its solubility.

The average atomic mass accounts for the different isotopes of an element and their natural abundances.

\[M_{\mathrm{avg.}}=\frac{M_{1}\times r_{1}+M_{2}\times r_{2}+M_{3}\times r_{3}}{r_{1}+r_{2}+r_{3}}\]

where M₁, M₂, M₃ are atomic masses of isotopes and r₁, r₂, r₃ are their relative abundances.

Percentage of an element in a compound \[=\frac{\text{Total wt. of the element in one molecule}}{\text{Gram molecular weight of the compound}}\times100\]

"In any chemical reaction, the total mass of reactants equals the total mass of products." Mass is neither created nor destroyed.

• Proposed by Antoine Lavoisier in 1789.
• e.g. Carbon in coal becomes CO₂ when burnt; the mass of C + O₂ = mass of CO₂.

"A given compound always contains exactly the same proportion of elements by mass, regardless of its source."

• Proposed by Joseph Proust in 1797.
• e.g. Pure water always has H : O mass ratio = 1 : 8, regardless of its source.
• e.g. Cupric carbonate (CuCO₃) found naturally or prepared synthetically always has the same percentage of Cu, C, and O.

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

• Proposed by Dalton in 1803.
• e.g. Hydrogen + Oxygen forms water (H₂O) and hydrogen peroxide (H₂O₂). With 2 g of H, oxygen combines as 16 g and 32 g → ratio = 1:2.

"Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules."

• Proposed by Avogadro in 1811.
• 1 mole of any gas at STP = 22.4 L (at 0°C, 1 atm) or 22.71 L (at 0°C, 1 bar — new IUPAC STP).
• 1 mole of any substance = 6.022 × 10²³ particles.

Avogadro's Law (Volume–Moles Relationship):

At constant temperature (T) and pressure (P), volume is directly proportional to number of moles.

Chemistry is central to all natural sciences and has applications in:

• Medicine and pharmaceuticals (drug design, antibiotics)
• Agriculture (fertilisers, pesticides)
• Industry (polymers, fuels, dyes)
• Environmental science (pollution control)
• Materials science (semiconductors, nanomaterials)

The SI system has 7 base units:

Temperature Conversions:

K = ^°C + 273.15

\[°F=\frac{9}{5}°C+32\]

Dalton's atomic theory laid the foundation of modern chemistry with four core postulates:

1. All matter is made up of extremely small particles called atoms.
2. Atoms of the same element are identical to each other in mass and properties; atoms of different elements differ.
3. Atoms can neither be created nor destroyed — they are indestructible.
4. Atoms combine in fixed, simple whole-number ratios to form compound atoms (molecules).

Note: Modern discoveries have refined some postulates (e.g., isotopes show atoms of the same element can differ in mass), but the core framework remains foundational.

Stoichiometry = calculation of relative masses of reactants and products involved in a chemical reaction, based on the balanced equation.

Steps to solve stoichiometric problems:

1. Balance the chemical equation.
2. Convert all given masses to moles.
3. Use mole ratios from the balanced equation to find moles of the wanted substance.
4. Convert moles back to the required unit (grams, litres, molecules).

Example: For 4Fe(s) + 3O₂(g) → 2Fe₂O₃(g)

4 mol Fe reacts with 3 mol O₂ to produce 2 mol Fe₂O₃.