The atomic structure and behavior of subatomic particles form the foundation of our understanding of matter and energy at the microscopic level.
The Rutherford Scattering experiment revolutionized our understanding of atomic structure by revealing that atoms have a dense, positively-charged nucleus surrounded by electrons. When alpha particles were fired at a thin gold foil, most passed straight through, but some bounced back at large angles. This unexpected observation led to the discovery that most of an atom's mass is concentrated in a tiny central nucleus, with electrons orbiting in the mostly empty space around it.
The mass difference between subatomic particles is crucial for understanding atomic stability. Protons and neutrons have nearly identical masses (about 1.67 × 10^-27 kg), though neutrons are slightly heavier. This small mass difference plays a vital role in nuclear stability and radioactive decay. Ionisation energy trends across the periodic table are influenced by several key factors: nuclear charge, electron shielding, and atomic radius. As you move left to right across a period, ionisation energy generally increases because the nuclear charge increases while the shielding effect remains relatively constant. Moving down a group, ionisation energy typically decreases as the outer electrons are farther from the nucleus and experience more shielding from inner electron shells. These trends help explain chemical reactivity and bonding patterns between elements.
The interplay between these atomic properties determines how elements behave in chemical reactions and combine to form compounds. Understanding these fundamental concepts helps explain everything from why noble gases are unreactive to why metals readily form positive ions. This knowledge forms the basis for modern chemistry and continues to guide research in fields ranging from materials science to nuclear physics.
