1.
Atoms and Their Composition:
o All matter is composed of tiny particles called atoms. These atoms serve as
the fundamental building blocks.
o Each atom consists of three subatomic particles:
Protons: Positively charged particles located in the nucleus.
Neutrons: Neutral particles also found in the nucleus.
Electrons: Negatively charged particles that orbit the nucleus in
specific energy levels called shells.
2. Relative Atomic Mass:
o Atoms are incredibly small, so we use a unit called relative atomic mass to
compare their masses.
o One relative atomic mass unit is equal to 1/12th the mass of a carbon-12 atom.
o Hydrogen, for instance, has a relative atomic mass of 1, meaning that 12
hydrogen atoms have the same mass as 1 carbon atom.
3. Atomic Number (Proton Number):
o The atomic number (Z) represents the number of protons in an atom’s
nucleus.
o It also corresponds to the number of electrons in a neutral atom.
o The atomic number determines an element’s position on the Periodic Table.
4. Mass Number (Nucleon Number):
o The mass number is the total count of protons and neutrons in an atom’s
nucleus.
o Neutrons can be calculated by subtracting the atomic number from the mass
number.
o Protons and neutrons together are referred to as nucleons.
5. Atomic Notation:
o Atomic notation provides a concise representation of an element’s atomic
structure.
o On the Periodic Table, you’ll find both the atomic number and the relative
atomic mass (which can be used as the mass number).
o For example, carbon’s atomic notation is written as: C (symbol) with Z = 6
(atomic number) and A = 12 (mass number).
The Rutherford Gold Foil Experiment
While at the University of Manchester, Rutherford collaborated with Hans Geiger and Ernest
Marsden on the gold-foil experiment. Here’s how it unfolded:
1. Alpha Particles: Rutherford had previously discovered alpha particles. In this
experiment, they directed these particles at an ultra-thin sheet of gold foil.
2. Expectations: Based on J.J. Thomson’s plum pudding model, Rutherford initially
expected the alpha particles to pass through the foil effortlessly.
3. Surprising Observations:
o Some alpha particles did indeed pass through the foil without deviation.
o However, others veered off course, bouncing back or deflecting.
o Rutherford likened this to shooting a bullet at tissue and having it rebound.
4. Two Key Deductions:
o Empty Space: Most of the atom consists of empty space.
o Atomic Nucleus: There must be something small, dense, and positively
charged within the atom to repel the positively charged alpha particles.
The Legacy of Rutherford’s Experiment
Rutherford’s gold foil experiment revolutionized our understanding of atomic structure. It
revealed the existence of the nucleus, a tiny, positively charged core at the center of the
atom. This pivotal discovery paved the way for subsequent atomic models, including Niels
Bohr’s model.
Classification and Arrangement:
o The Periodic Table is a systematic arrangement of chemical elements based on
their atomic number (number of protons).
o Elements are grouped into periods (rows) and groups (columns).
o Periods represent the energy levels (shells) of electrons, while groups share
similar chemical properties due to their similar electron configurations.
2. Periods and Blocks:
o The Periodic Table has 7 periods (numbered 1 to 7).
o Elements within the same period have the same number of electron shells.
o The blocks are named after the type of subshell being filled:
s-block: Groups 1 and 2 (alkali metals and alkaline earth metals).
p-block: Groups 13 to 18 (nonmetals, metalloids, and noble gases).
d-block: Transition metals (Groups 3 to 12).
f-block: Lanthanides and actinides (rare earth elements).
3. Groups and Trends:
o Group 1 (Alkali Metals):
Highly reactive metals.
Form +1 ions.
React vigorously with water.
o Group 17 (Halogens):
Nonmetals.
Form -1 ions.
Exist as diatomic molecules (e.g., Cl₂, Br₂).
o Group 18 (Noble Gases):
Inert gases.
Stable electron configurations.
Rarely react with other elements.
4. Atomic and Ionic Radii:
o Atomic radius decreases across a period due to increased nuclear charge.
o Atomic radius increases down a group due to additional electron shells.
o Ionic radius follows similar trends.
5. Ionization Energies:
o First ionization energy: Energy required to remove one mole of electrons from
one mole of atoms.
o Generally increases across a period and decreases down a group.
o Exceptions occur due to electron pairing and subshell stability.
6. Electronegativity:
o Ability of an atom to attract electrons in a chemical bond.
o Increases across a period and decreases down a group.
o Fluorine is the most electronegative element.
7. Metallic and Non-metallic Properties:
o Metals: Good conductors, malleable, ductile.
o Nonmetals: Poor conductors, brittle, often form covalent compounds.
