Revision: Some Basic Concepts in Chemistry JEE Main Some Basic Concepts in Chemistry

An atom is the smallest particle of a chemical element that retains its chemical properties.

Metal:  A chemical element that is an effective conductor of electricity and heat can be defined as a metal.

Ex.: Copper, Iron, Silver, etc.

Metalloid: Metalloid is a chemical element that exhibits some properties of metals and some of non-metals. Metalloids are generally semi-conductors.

Ex.: Silicon. Arsenic, Antimony and Boron.

An atom is the smallest particle of an element which retains its chemical identity in all physical and chemical changes.

Element: It is a substance that cannot be broken down into simpler substance by chemical means

Ex.: Oxygen, Hydrogen, Gold & Helium.

An Atom: Smallest particle of an element that can exist and have properties of an element.

Atom: An atom is the smallest indivisible unit of an
OR
Atom is the smallest unit of matter.

Relative atomic mass— Relative atomic mass is the mass of an atom of an element as a multiple of the standard atomic mass unit.

The relative atomic mass of an element is the ratio between the average mass of its isotopes to 1/12th part of the mass of a carbon – 12 atoms. It is denoted as A_r.

Relative atomic mass = `" Average mass of the isotopes of the element"/(1"/"12^{"th"}" of the mass of one Carbon- 12 atom")`

Compound: A compound is a pure substance that is formed when the atoms of two or more elements combine chemically in definite proportions.

Ex: H20, NaCl.

Non-Metal: Non-metal is an element that doesn’t have the characteristics of metal including, (i.e.) ability to conduct heat or electricity luster or flexibility.

Ex. Carbon Iodine, Sulphur.

Mass number— Mass number is the sum of the number of protons and neutrons present in the nucleus of an atom. It is denoted by A.

Radicals : A radical is an atom of an element or a group of atoms of different elements that behaves as a single unit with a positive or negative charge on it.

An atom which becomes charged by losing or gaining electrons is called an ion.

Molecule : Molecule is the smallest unit of a compound (or an element) which always has an independent existance.

Covalent bond— When atoms of different non-metals neither donate nor accept electrons and hence no ions are formed, such a bond is called covalent bond.

Chemical bond— A chemical bond is the binding force between two or more atoms of a molecule.

“Mixtures can be defined as. a kind of matter which is formed by mixing two or more pure substances (elements and compounds) in any proportion, such that they do not undergo any chemical change and retain their individual properties. Therefore they are impure substances.

Element is a substance which cannot be broken further into simpler substances and has a definite set of properties. Elements are made up of only one kind of atoms.

Formula: Formula is a short way of representing the molecule of an element or a compound

Atom: An atom is the smallest indivisible unit of an element which exhibits all the properties of that element and may or may not have an independent existence. An atom is the smallest indivisible unit of an element which exhibits all the properties of that element and may or may not have an independent existence. 

Molecule: A molecule can be defined as the smallest unit of an element or a compound which exhibits all the properties of that element or compound and has an independent existence. They are divisible into atoms.

The number of atoms in a molecule of an element is called its atomicity. 

Compounds are pure substances composed of two or more elements in definite proportion by mass and has properties, entirely different from those of its constituents elements.
Compound, are made up of different types of atoms combined chemically.

Element is a substance which cannot be split up into two or more simple substances by usual chemical methods of applying heat, light or electric energy; for example, hydrogen, oxygen and chlorine.

Elements: An element is defined as a pure substance made up of only one kind of atoms that cannot be converted into anything simpler than itself by any physical or chemical process. 

Compounds: Compounds are pure substances composed of two or more elements in definite proportion by mass and has a definite set of properties. The compound is made up of only one kind of molecules

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.

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.

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

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

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 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.

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

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.

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

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.

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

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

• 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.

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

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

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 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.

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.

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

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.

"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.

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

\[°F=\frac{9}{5}°C+32\]

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.