Atomic Structure and Isotopes

Atomic Structure & Subatomic Particles

• Atoms consist of protons, neutrons, and electrons.
• Protons (p⁺) possess a positive charge and reside in the nucleus; the quantity of protons defines the atomic number (Z) of an element.
• Neutrons (n⁰) are uncharged particles located in the nucleus, contributing to the mass number (A).
• Electrons (e⁻) have a negative charge and move around the nucleus in electron clouds (orbitals); electrons are much smaller than protons and neutrons.
• Isotopes arise from variations in neutron number within atoms of the same element.
• Neutrons = Mass Number - Atomic Number

Isotopes & Radioisotopes

• Isotopes are variants of an element with the same proton number but differing neutron numbers, thus having different mass numbers.
• Carbon-12 (C-12) contains 6 protons and 6 neutrons.
• Carbon-14 (C-14) contains 6 protons and 8 neutrons.
• Radioisotopes are unstable isotopes that emit radiation over time.
• Carbon-14 serves in radiocarbon dating for determining fossil ages.
• Uranium-235 is utilized in nuclear power plants and atomic bombs.
• Iodine-131 is employed in treating thyroid cancer.
• Cobalt-60 is utilized in radiotherapy for cancer treatment.
• Radioactive decay involves an unstable nucleus emitting alpha (α), beta (β), or gamma (γ) rays.

Atomic Models & Their Creators

• John Dalton (1803) proposed the Billiard Ball Model: atoms are solid, indivisible spheres.
• J.J. Thomson (1897) discovered the electron and proposed the Plum Pudding Model: electrons scattered within a positively charged "pudding."
• Ernest Rutherford (1911) performed the Gold Foil Experiment, finding that atoms have a small, dense, positively charged nucleus surrounded by empty space and proposed the Nuclear Model.
• Niels Bohr (1913) refined the model to the Bohr Model: electrons orbit the nucleus in fixed energy levels (shells).
• Erwin Schrödinger (1926) introduced the Quantum Model: electrons exist in probability clouds (orbitals) rather than fixed paths.
• James Chadwick (1932) discovered the neutron, explaining isotopes (different masses of the same element).
• Rutherford's Gold Foil Experiment demonstrated that atoms are mostly empty space with a small, dense nucleus, as some alpha particles were deflected.

Periodic Table Overview

• The periodic table orders elements by atomic number and groups them by chemical properties.
• Elements in the same group (column) share similar valence electron numbers, determining their reactivity.
• Alkali Metals (Group 1: Li, Na, K, Rb, Cs, Fr) are extremely reactive with 1 valence electron, react violently with water, and have soft textures and low melting points.
• Alkaline Earth Metals (Group 2: Be, Mg, Ca, Sr, Ba, Ra) are less reactive than Group 1, possess 2 valence electrons, and are found in bones and fireworks.
• Transition Metals (Groups 3-12: Fe, Cu, Ag, Au, etc.) are hard, dense, have high melting points and good conductivity, form colored compounds, and exhibit multiple oxidation states.
• Halogens (Group 17: F, Cl, Br, I, At) are highly reactive non-metals with 7 valence electrons, form salts with metals (e.g., NaCl), and exist in various states at room temperature.
• Noble Gases (Group 18: He, Ne, Ar, Kr, Xe, Rn) are unreactive due to full outer electron shells and are used in neon signs and balloons.
• Metalloids (B, Si, Ge, As, Sb, Te) are found along the staircase and have properties of both metals and non-metals.
• Silicon (Si) is a key semiconductor element.

Periodic Table Trends

• Atomic Radius (Size) increases down a group due to added electron shells.
• Atomic Radius (Size) decreases across a period because the nucleus pulls electrons closer.
• Electronegativity (Ability to attract electrons) increases across a period (except noble gases).
• Electronegativity (Ability to attract electrons) decreases down a group as atoms get larger.
• Ionization Energy (Energy to remove an electron) increases across a period as nuclear attraction strengthens.
• Ionization Energy (Energy to remove an electron) decreases down a group as electrons are farther from the nucleus.

Common Elements & Their Uses

• Hydrogen (H) is the most abundant element, used in fuel and water.
• Oxygen (O) is essential for respiration and combustion.
• Carbon (C) is the basis of organic life and is found in diamonds and coal.
• Iron (Fe) is used in steel production and blood hemoglobin.
• Copper (Cu) is used in electrical wiring due to its conductivity.
• Gold (Au) and Silver (Ag) are used in jewelry and electronics.
• Uranium (U) is used in nuclear energy production.