Atomic Structure
Section: 2. Atoms, Elements & Compounds | Syllabus: Cambridge AS Level Physics 9702
The Atom Atoms are the basic building blocks of all matter. Everything around us is made of atoms, and understanding their structure is fundamental to chemistry. Atom The smallest particle of an element that retains the chemical properties of that element.
Atoms cannot be divided by chemical means. Subatomic Particles Atoms are made up of three types of subatomic particles: protons, neutrons, and electrons. Particle Location Relative Mass Relative Charge Actual Mass (kg) Proton Nucleus 1 +1 1.673 × 10⁻²⁷ Neutron Nucleus 1 0 (neutral) 1.675 × 10⁻²⁷ Electron Shells/Orbitals around nucleus 1/1840 (negligible) -1 9.109 × 10⁻³¹ Key Point Protons and neutrons have approximately the same mass.
Electrons are about 1/1840th the mass of a proton, so their mass is considered negligible when calculating atomic mass. Structure of the Atom The Nucleus Located at the center of the atom Contains protons and neutrons (collectively called nucleons) Very small compared to the atom - about 1/10,000th the diameter Contains almost all the atom's mass Has a positive charge due to protons Electron Shells Electrons orbit the nucleus in shells (also called energy levels) Shells are at different distances from the nucleus Electrons in shells further from the nucleus have higher energy The region where electrons orbit is mostly empty space Amazing Fact Atoms are mostly empty space!
If an atom were the size of a football stadium, the nucleus would be about the size of a pea at the center, with electrons orbiting in the stands. Atomic Number and Mass Number Atomic Number (Proton Number) Atomic Number (Z) The number of protons in the nucleus of an atom.
This defines what element the atom is. Symbol: Z Unique to each element Determines the element's identity (e.g., all carbon atoms have 6 protons) In a neutral atom, atomic number = number of electrons Found on the periodic table (smaller number) Mass Number (Nucleon Number) Mass Number (A) The total number of protons and neutrons in the nucleus of an atom.
Symbol: A Mass number = protons + neutrons Can vary for the same element (isotopes) Found on the periodic table (larger number, often rounded) Number of neutrons = Mass number - Atomic number Neutrons = A - Z Notation for Atoms Atoms are represented using standard notation: Mass number ₍Atomic number₎ Element Symbol Example: 12 ₆C or 23 ₁₁Na Mass number (A) is written as a superscript (top left) Atomic number (Z) is written as a subscript (bottom left) Element symbol is in the middle Example 23 ₁₁Na represents a sodium atom with 11 protons, 12 neutrons (23-11=12), and 11 electrons.
Calculating Subatomic Particles Example 1: Carbon-12 ( 12 ₆C) Protons: 6 (atomic number) Neutrons: 12 - 6 = 6 Electrons: 6 (same as protons in neutral atom) Example 2: Chlorine-35 ( 35 ₁₇Cl) Protons: 17 (atomic number) Neutrons: 35 - 17 = 18 Electrons: 17 (same as protons in neutral atom) Why Atoms are Neutral Atoms have no overall electrical charge because: Number of protons (+) = Number of electrons (-) Positive and negative charges cancel out Example: Carbon has 6 protons (+6) and 6 electrons (-6), total charge = 0 Important If an atom gains or loses electrons, it becomes charged and is called an ion.
Gaining electrons makes it negatively charged (anion), losing electrons makes it positively charged (cation). Relative Atomic Mass Because atoms are so tiny, we use relative atomic mass instead of actual mass.
Relative Atomic Mass (Ar) The average mass of atoms of an element compared to 1/12th the mass of a carbon-12 atom. It accounts for the different isotopes and their abundances. Based on the carbon-12 standard (assigned mass of exactly 12) Has no units (it's a ratio) Usually close to the mass number but may not be a whole number Takes into account isotopes and their relative abundances Example Chlorine has a relative atomic mass of 35.5 (not a whole number) because it exists as a mixture of isotopes: Cl-35 (75%) and Cl-37 (25%).
The weighted average gives 35.5. Development of Atomic Models Our understanding of atomic structure has evolved over time: 1. Dalton's Model (1803) Atoms are tiny, indivisible spheres Different elements have different types of atoms First scientific atomic theory 2.
Thomson's Plum Pudding Model (1897) Atoms contain negative electrons Electrons embedded in a positive "pudding" Discovered after Thomson found electrons 3. Rutherford's Nuclear Model (1911) Atoms have a small, dense, positive nucleus Electrons orbit the nucleus Most of the atom is empty space Based on the famous gold foil experiment 4.
Bohr's Model (1913) Electrons orbit in specific shells/energy levels Electrons can jump between shells by absorbing/emitting energy Each shell can hold a maximum number of electrons 5. Modern Quantum Model (Current) Electrons exist in orbitals (regions of probability) Cannot determine exact position and momentum simultaneously Based on quantum mechanics Key Experiment Rutherford's gold foil experiment: Alpha particles were fired at th…
Interactive revision notes, videos and practice questions load below.