Revision: Std. XI >> Some Basic Concepts of Chemistry MAH-MHT CET (PCM/PCB) Some Basic Concepts of Chemistry

Chemistry is the branch of science that deals with the identification of substances, the matter they are composed of, and the investigation of their properties.

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.

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. 

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

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 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.

One mole is the amount of a substance that contains as many entities or particles as there are atoms in exactly 12 g of the carbon-12 isotope.

or

One mole is the amount of substance which contains 6.022 ×10²³ (avogadro's number) particles/entities (such as atoms, molecules or ions).

e.g.

• 1 mole of nitrogen atoms = 6.022 ×10²³ atom of nitrogen
• 1 mole of water molecules = 6.022 ×10²³ molecule of water. 
• 1 mole of sodium bromide = 6.022 ×10²³ formula unit of NaBr.

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.

Mass % of a Component (w/w) \[=\frac{\text{Mass of the component in the solution}}{\text{Total mass of the component}}\times100\]

\[\mathrm{Molarity~}(M)=\frac{\text{Number of moles of solute}}{\text{Volume of solution (in L)}}\]

\[\mathrm{Normality~}(N)=\frac{\text{Number of gram equivalents}}{\text{Volume of solution (in L)}}\]

\[\mathrm{Molality~}(m)=\frac{\text{Moles of solute}(n)}{\text{Mass of solvent}(W_A)\mathrm{~in~kg}}\]

\[\mathrm{n=\frac{Mass~of~a~substance}{Molar~mass~of~a~substance}}\]

\[\mathrm{Number of Moles of Gas=\frac{Volume~of~the~gas~at~STP}{Molar~mass~of~a~substance}}\]

Five fundamental laws govern how elements and compounds combine chemically:

Law 1 — Law of Conservation of Mass (Antoine Lavoisier)

Mass is neither created nor destroyed during any chemical reaction. The total mass of reactants always equals the total mass of products.

Law 2 — Law of Definite Proportion (Joseph Proust)

A specific chemical compound always contains its elements combined in a fixed ratio by weight, regardless of where the compound comes from or how it was made.

Exception: This law does not hold for compounds made from different isotopes of an element.

Law 3 — Law of Multiple Proportion (John Dalton)

When two elements combine to form more than one compound, the different masses of one element that combine with a fixed mass of the other are always in a simple whole-number ratio. Example: CO and CO₂.

Law 4 — Gay Lussac's Law of Gaseous Volumes

When gases react or are produced in a chemical reaction, their volumes bear a simple whole-number ratio to each other — provided temperature and pressure remain the same.

Law 5 — Avogadro's Law

At the same temperature and pressure, equal volumes of all gases contain the same number of molecules, regardless of the type of gas.

"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.

Matter is categorised based on its chemical composition into two broad groups:

1. Pure Substances have a definite, fixed chemical composition. They are further divided into:

• Elements — the simplest form of matter; cannot be broken down further by ordinary chemical means. Example: pure silver.

• Compounds — formed when two or more elements chemically combine in a fixed ratio. Example: common salt (NaCl).

2. Mixtures have no fixed composition and therefore no definite properties. They are divided into:

• Homogeneous Mixtures — constituents are uniformly distributed throughout the sample. Example: vinegar.

• Heterogeneous Mixtures — constituents are not uniformly distributed. Example: tomato sauce.

Quick memory trick:

Pure → Fixed composition. Mixture → Variable composition.

Types of Properties

• Physical Properties — can be observed or measured without altering the chemical nature of the substance. Examples: colour, odour, melting point, boiling point, density.

• Chemical Properties — involve a chemical change in the substance; the original substance is converted into something new. Example: burning coal produces CO₂.

SI Fundamental Units

The International System of Units (SI) defines seven base units that serve as building blocks for all scientific measurement:

Key Notes to Remember:

• Mass measures the quantity of matter and is independent of location. Weight depends on gravity — the same object has different weight on Earth vs. the Moon, but identical mass.
• Temperature and heat are not the same. Heat is energy being transferred; temperature tells us the direction of that transfer.
• 0°C = 32°F; 100°C = 212°F. A rise of 1°C corresponds to a rise of 9/5°F on the Fahrenheit scale.
• Units can be written in two equivalent ways: g/cm³ or g cm⁻³ — both are acceptable.

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.

• The molar mass of any element in grams is numerically equal to its atomic mass in u.
• For polyatomic molecules, molar mass in grams equals the molecular or formula mass in u.
• "1 mole oxygen atoms" and "1 mole oxygen molecules" are NOT the same — always specify what entities you are counting.
• Earlier, the unit amu was used; it has been replaced by u (unified atomic mass unit).