Revision: Class 11 >> Some Basic Concepts of Chemistry NEET (UG) Some Basic Concepts of Chemistry

Chemistry is defined as the branch of science that deals with the study of the composition, structure, and properties of matter.

Anything that has mass and occupies space is called matter.

The process by which matter changes from one state to another and back to the original state, without any change in its chemical composition.

Unit is the quantity of a constant magnitude which is used to measure the magnitudes of other quantities of the same nature.

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.

Precision is about getting reproducible results. If you measure the same thing multiple times and get nearly identical answers, your measurements are precise.

or

The quality, condition, or fact of being exact and accurate or the closeness of the set of values obtained from identical measurements of quantity.

Precision = Individual Value - Arithmetic Mean Value

Accuracy is about how close your measured value is to the true, actual value of that quantity.

or

The quality or state cate of being accurate or the ability to work or perform without making mistakes.

Accuracy = Mean value - True Value

In real experiments, it is very difficult to get exactly the same answer every single time. This difference or possibility of error is called uncertainty.

The reactant which is completely used up in a reaction is known as Limiting reagent or Limiting reactant.

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.

The sum of the atomic masses of the atoms present in the formula is called the formula 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.

A formula that gives the simplest whole number ratio of atoms of each element in a compound.

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.

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.

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

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

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

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

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

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

Shows the exact number of atoms of each element present in one molecule of a compound.

Molecular formula = (Empirical formula)_n​

\[n=\frac{\text{Molecular mass}}{\text{Empirical formula mass}}\]

\[x_A=\frac{n_A}{n_A+n_B}\quad x_B=\frac{n_B}{n_A+n_B}\]

x_A​ + x_B ​= 1

Mass % of element\[=\frac{\text{Mass of element in compound}}{\text{Molar mass of compound}}\times100\]

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.

"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.

"The ratio in which two elements A and B separately combine with a fixed mass of a third element C is either the same as, or a simple multiple of, the ratio in which A and B combine directly with each other."

• Also called the Law of Equivalences.
• e.g. Carbon and sulfur both combine with oxygen; the ratio in which they combine with oxygen is related to the ratio in which they combine with each other.

"When gases combine or are produced in a chemical reaction, they do so in a simple ratio by volume, provided all gases are at the same temperature and pressure."

• Proposed by Gay-Lussac in 1808.
• e.g. 100 mL H₂ + 50 mL O₂ → 100 mL H₂O vapour (ratio = 2 : 1 : 2).
• The volume ratio of gaseous reactants to products agrees with their molar ratio.
• Volume of a gas is directly proportional to the number of moles (not inversely).

Statement: The volume remaining constant, the pressure of a given mass of gas increases or decreases by 1/273.15 of its pressure at 0°C for each 1°C rise or fall in temperature.

where β = pressure expansion coefficient = 1/273 per °C.

"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)

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.

The SI system has 7 base units:

Temperature Conversions:

K = ^°C + 273.15

\[°F=\frac{9}{5}°C+32\]

• Fundamental (base) units — independent units for fundamental quantities (length, mass, time, etc.)
• Derived units — combinations of base units (e.g., m/s for speed, kg/m³ for density, Pa = kg m⁻¹ s⁻² for pressure)

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).

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.