Revision: Atomic Structure Science SSLC (English Medium) Class 8 Tamil Nadu Board of Secondary Education

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 a chemical element that retains its chemical properties.

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

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

Ex.: Oxygen, Hydrogen, Gold & Helium.

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

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

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.

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

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

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.

Atomic number refers to the number of protons present in an atom. It is denoted by Z. Example: An atom of oxygen contains 8 proton Therefore its atomic number is 8.

Mass number refers to the sum of the number of protons and neutrons present in the nucleus of an atom and denoted by A Mass number = Number of protons + Number of neutrons.

The outermost shell of an atom is known as its valence shell.

The protons and neutrons collectively are known as nucleons.

The valency of an element is determined by the number of electrons present in the outermost shell of its atoms, that is, the valence electrons.

When the properties of elements in a period or a group of the modern periodic table are compared, certain regularity is observed in their variations. It is called the periodic trends in the modern periodic table.

The reactions in which heat is absorbed are called endothermic reactions. The reactants absorb heat for form products.

A chemical equation is a balanced account of a chemical transaction. It is not merely a qualitative statement, but it also gives quantitative information of a chemical reaction.

OR

The representation of a chemical reaction in a condensed form using chemical formulae is called as the chemical equation.

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.

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.

• Discovered by J. J. Thomson (1897) using the cathode ray tube experiment
• Charge = –1.6 × 10⁻¹⁹ C
• Mass = 9.109 × 10⁻³¹ kg (very small)
• Cathode rays travel in straight lines
• Deflected by electric and magnetic fields (proves negative charge)

• Discovered by E. Goldstein using a discharge tube (canal rays)
• Charge = +1.6 × 10⁻¹⁹ C
• Mass = 1.673 × 10⁻²⁷ kg
• Present in the nucleus of an atom
• Determines the atomic number (Z) of an element

• Discovered by James Chadwick (1932)
• Charge = 0 (neutral)
• Mass = 1.675 × 10⁻²⁷ kg (almost equal to proton)
• Present in the nucleus along with protons
• Responsible for isotopes and atomic mass

Reaction: \[_4^9Be+_2^4He\to_6^{12}C+_0^1n\]

• Proposed by J. J. Thomson in 1904 after the discovery of electrons.
• The atom is a uniform sphere of positive charge.
• Electrons are embedded within this sphere.
• The positive charge is spread evenly throughout the atom.
• Total positive charge = total negative charge, so the atom is neutral.
• The model explained the presence of electrons in atoms.
• It did not include a nucleus in the atom.
• It failed to explain Rutherford’s results from the gold foil experiment.

• Word equations use names; chemical equations use formulas.
• Reactants → Products, with arrow showing reaction direction.
• Use + between two or more reactants or products.
• Show states: (s), (l), (g), (aq); use ↑ for gas, ↓ for precipitate.
• Heat (Δ) or other conditions go above/below the arrow.

• Law of Conservation of Mass: In a chemical reaction, mass is neither created nor destroyed, so the number of atoms of each element must be equal on both sides.
• A skeletal (unbalanced) equation has unequal atoms of one or more elements on the LHS and RHS.
• Balancing is done using the hit-and-trial method, starting with the compound having the most atoms and balancing hydrogen and oxygen last.
• Only coefficients are changed while balancing; chemical formulas must not be altered.
• A balanced equation may also indicate physical states (s, l, g, aq) and reaction conditions, such as temperature, pressure, or a catalyst.