Complete Chemistry notes on Atomic Structure for BPSC and Other Competitive Exams in 2025
General Science- 8 (Chemistry)
Understanding Atomic Structure and the basics of the Periodic Table is fundamental for the BPSC CCE General Science section.
1. Atomic Structure
An atom is the smallest unit of an element that retains the chemical identity of that element. It’s composed of subatomic particles:
- Protons (p+):
- Location: Found in the nucleus (the central part of the atom).
- Charge: Positive (+1 elementary charge).
- Mass: Approximately 1 atomic mass unit (amu).
- Significance: The number of protons defines the element. This is known as the Atomic Number (Z).
- Neutrons (n0):
- Location: Found in the nucleus, along with protons.
- Charge: Neutral (0 charge).
- Mass: Approximately 1 atomic mass unit (amu), very slightly heavier than a proton.
- Significance: Contribute to the atom’s mass and nuclear stability. Atoms of the same element can have different numbers of neutrons, leading to isotopes.
- Electrons (e−):
- Location: Orbit the nucleus in specific energy levels or shells.
- Charge: Negative (-1 elementary charge).
- Mass: Extremely small, approximately 18361 of a proton’s mass (negligible in comparison to protons and neutrons).
- Significance: Determine the chemical properties and reactivity of an atom. In a neutral atom, the number of electrons equals the number of protons.
Key Definitions:
- Atomic Number (Z): The number of protons in the nucleus of an atom. It uniquely identifies an element. For a neutral atom, Atomic Number = Number of Protons = Number of Electrons.
- Example: All carbon atoms have 6 protons (Z=6).
- Mass Number (A): The total number of protons and neutrons in the nucleus of an atom.
- Mass Number (A) = Number of Protons (Z) + Number of Neutrons.
- Number of Neutrons = Mass Number (A) – Atomic Number (Z).
- Example: Carbon-12 has a mass number of 12 (6 protons + 6 neutrons).
- Isotopes: Atoms of the same element (same atomic number, Z) that have different numbers of neutrons (and thus different mass numbers, A). They have identical chemical properties but different physical properties (e.g., density).
- Examples: Hydrogen has three isotopes: Protium (11H), Deuterium (12H), and Tritium (13H). Carbon has Carbon-12 (612C) and Carbon-14 (614C).
Brief Overview of Atomic Models:
- Dalton’s Atomic Theory (early 19th century): Atoms are indivisible, indestructible particles; all atoms of an element are identical. (Partially disproved by subatomic particles and isotopes).
- J.J. Thomson’s Plum Pudding Model (1897): Atom is a sphere of positive charge with electrons embedded in it. (Disproved by Rutherford).
- Rutherford’s Nuclear Model (1911): Atom has a small, dense, positively charged nucleus with electrons orbiting around it. (Proposed the existence of the nucleus).
- Niels Bohr’s Model (1913): Electrons orbit the nucleus in specific, quantized energy levels or shells. Electrons can jump between these levels by absorbing or emitting energy. (Introduced the concept of energy shells, which is still relevant).
- Quantum Mechanical Model (modern): Describes electrons in terms of probabilities (electron clouds or orbitals) rather than fixed orbits. This is the most accurate model, but for CCE, understanding Bohr’s concept of energy shells is often sufficient.
2. Periodic Table Basics
The Periodic Table is a tabular arrangement of chemical elements, ordered by their atomic number, electron configurations, and recurring chemical properties. It’s a powerful tool for understanding chemical behavior.
- Periods (Rows):
- There are 7 periods (numbered 1 to 7).
- Elements in the same period have the same number of electron shells (e.g., all elements in Period 3 have 3 electron shells).
- As you move from left to right across a period, atomic number increases, and properties change gradually from metallic to non-metallic.
- Groups/Families (Columns):
- There are 18 groups (numbered 1 to 18).
- Elements within the same group generally have the same number of valence electrons (electrons in the outermost shell), which determines their chemical properties.
- Therefore, elements in the same group exhibit similar chemical properties.
Key Groups and Their Properties:
- Group 1: Alkali Metals (Li, Na, K, etc.): Highly reactive metals, form +1 ions, soft, low density, react vigorously with water.
- Group 2: Alkaline Earth Metals (Be, Mg, Ca, etc.): Reactive metals, form +2 ions, slightly harder than alkali metals.
- Groups 3-12: Transition Metals: Hard, high melting/boiling points, good conductors, often form colored compounds, can exhibit variable valency.
- Group 17: Halogens (F, Cl, Br, I, etc.): Highly reactive non-metals, form -1 ions, exist as diatomic molecules (F2, Cl2).
- Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn): Extremely unreactive (inert) due to a full outermost electron shell (stable octet/duet), exist as monatomic gases.
Metals, Non-metals, and Metalloids:
- Metals: (Left and center of the table)
- Properties: Lustrous (shiny), malleable (can be hammered into sheets), ductile (can be drawn into wires), good conductors of heat and electricity, usually solid at room temperature (except Mercury). Tend to lose electrons to form positive ions.
- Non-metals: (Upper right side of the table)
- Properties: Dull, brittle (if solid), poor conductors of heat and electricity (insulators), exist as solids, liquids, or gases at room temperature. Tend to gain electrons to form negative ions.
- Metalloids: (Along the zigzag line between metals and non-metals: B, Si, Ge, As, Sb, Te, Po)
- Properties: Exhibit properties intermediate between metals and non-metals (e.g., semiconductors).
Valence Electrons: These are the electrons in the outermost occupied electron shell of an atom. They are primarily responsible for an atom’s chemical behavior and its ability to form bonds.
Periodic Trends (General Tendencies):
- Atomic Size (Atomic Radius):
- Decreases across a period (left to right): Due to increasing nuclear charge pulling electrons closer.
- Increases down a group (top to bottom): Due to the addition of new electron shells.
- Ionization Energy: The energy required to remove an electron from a gaseous atom.
- Increases across a period (left to right): Due to stronger nuclear pull on valence electrons.
- Decreases down a group (top to bottom): Due to increasing distance of valence electrons from the nucleus and increased shielding.
- Electronegativity: The ability of an atom to attract electrons in a chemical bond.
- Increases across a period (left to right): Due to increasing nuclear charge.
- Decreases down a group (top to bottom): Due to increasing atomic size and shielding.
- Note: Fluorine (F) is the most electronegative element.
Objective Questions
Here are 10 objective questions related to Atomic Structure and Periodic Table basics, keeping the BPSC CCE format in mind:
Atomic Structure and Periodic Table Basics – Questions for Revision
- Which subatomic particle has a positive charge and is found in the nucleus of an atom?
- Correct Answer: Proton
- The atomic number (Z) of an element is determined by the number of:
- Correct Answer: Protons
- Atoms of the same element that have different numbers of neutrons are called:
- Correct Answer: Isotopes
- Which of the following elements is an Alkali Metal?
- Correct Answer: Sodium (Na)
- As you move from left to right across a period in the Periodic Table, what generally happens to the atomic size?
- Correct Answer: It decreases.
- Which group of elements is known for being largely unreactive or ‘inert’?
- Correct Answer: Noble Gases
- The outermost electrons that determine an atom’s chemical properties are called:
- Correct Answer: Valence electrons
- An element with atomic number 17 and mass number 35 has how many neutrons?
- Correct Answer: 18
- Elements in the same group of the Periodic Table generally have similar:
- Correct Answer: Chemical properties
- Which of the following is true about non-metals?
- Correct Answer: They are generally dull and brittle (if solid).
