Matter:

• Anything that has mass and takes up space.

Models:

• Simplified representations of complex physical phenomena.

• Used to explain and predict physical behaviors.

Ancient Greek Theories

• Democritus (circa 460–370 BC):

  • Proposed that matter is composed of small, indivisible particles called "atomos," meaning "uncuttable."

  • Believed that atoms were different shapes and sizes, which determined the properties of the matter they composed.

  • His ideas were purely philosophical and lacked experimental evidence.

• Aristotle (384–322 BC):

  • Rejected Democritus's idea of atoms.

  • Suggested that matter is continuous and composed of four fundamental elements: earth, water, air, and fire.

  • Introduced the concept of "hylomorphism," which posits that all things are a combination of matter (hyle) and form (morphe).

John Dalton (1766–1844)

• Dalton's Atomic Theory (1803):

  • Proposed that matter is composed of small, indivisible particles called atoms.

  • Each element consists of identical atoms, which differ from those of other elements.

  • Compounds are formed by the combination of atoms of different elements in fixed ratios.

  • Chemical reactions involve the rearrangement of atoms; atoms are not created or destroyed in these reactions.

• Significance:

  • Provided the first scientific basis for the concept of the atom.

  • Dalton’s work laid the groundwork for modern chemistry and the law of multiple proportions.

J.J. Thomson (1856–1940)

• Discovery of the Electron (1897):

  • Conducted experiments using cathode ray tubes.

  • Demonstrated that cathode rays were composed of negatively charged particles, later named electrons.

• Plum Pudding Model (1904):

  • Proposed that the atom is a sphere of positive charge with electrons embedded within it, like raisins in a “plum pudding”.

  • Suggested that the positive and negative charges were distributed evenly throughout the atom.

Ernest Rutherford (1871–1937)

• Gold Foil Experiment (1911):

  • Conducted by Hans Geiger and Ernest Marsden under Rutherford’s supervision.

  • Involved directing alpha particles at a thin gold foil.

  • Most alpha particles passed through the foil, but some were deflected at large angles.

• Nuclear Model (1911):

  • Concluded that atoms have a small, dense, positively charged nucleus at the center.

  • Electrons orbit the nucleus at a relatively large distance.

  • Most of the atom’s volume is empty space.

• Significance:

  • Overturned the plum pudding model.

  • Provided the basis for understanding atomic structure and led to the development of nuclear physics.

Niels Bohr (1885–1962)

• Bohr Model (1913):

  • Proposed that electrons orbit the nucleus in fixed, quantized orbits or energy levels.

  • Electrons can jump between orbits by absorbing or emitting specific amounts of energy (quantum).

  • Introduced the concept of quantized angular momentum.

• Significance:

  • Explained the spectral lines of hydrogen.

  • Incorporated quantum theory into atomic structure.

  • Laid the foundation for quantum mechanics.

Modern Atomic Model

• Quantum Mechanical Model:

  • Developed through the work of scientists like Schrödinger, Heisenberg, and Dirac.

  • Describes electrons in terms of probability distributions rather than fixed orbits.

  • Schrödinger Equation: Describes how the quantum state of a physical system changes over time.

  • Heisenberg Uncertainty Principle: States that it is impossible to simultaneously know the exact position and momentum of an electron.

  • Orbitals: Regions in space where the probability of finding an electron is high.

• Significance:

  • Provides a more accurate and comprehensive understanding of atomic and subatomic particles.

  • Essential for the development of technologies like semiconductors, lasers, and nuclear energy.

States of matter

Solids

• Fixed, definite shape and volume

• Very strong intermolecular forces

• Particles closely packed

• Particles vibrate about their mean position

• No fluidity

• Incompressible

• Particles arranged in fixed positions and in a regular pattern

Liquids

• Indefinite shape and definite volume

• Strong intermolecular forces

• Particles are close but not packed

• Particles can slide over each other

• Can flow

• Incompressible

• Expand more than solids if heated

Gases

• Indefinite shape and volume

• Can flow

• Particles have a large gap between them

• Very weak intermolecular forces

• Particles far apart from each other

• Compressible

• Particles are in continuous, random motion

• Expand more than liquids and gases if heated

Phase Transitions

• Melting: Solid to liquid.

• Freezing: Liquid to solid.

• Vaporization: Liquid to gas (includes boiling and evaporation).

• Condensation: Gas to liquid.

• Sublimation: Solid to gas without passing through the liquid phase.

• Deposition: Gas to solid without becoming liquid.

Kinetic Theory of Gases

• Assumptions:

  • Gas particles are in constant, random motion.

  • Collisions between gas particles are perfectly elastic.

  • The volume of gas particles is negligible compared to the volume of their container.

  • There are no intermolecular forces acting between gas particles.

Ideal Gas Law

Thermal Properties of Matter

• Temperature: Measure of the average kinetic energy of particles.

• Heat: Transfer of thermal energy between substances due to a temperature difference.

• Specific Heat Capacity: Amount of heat required to change the temperature of 1 kg of a substance by 1°C.

• Latent Heat: Heat required for a phase change without a change in temperature.

  • Latent Heat of Fusion: Heat required to change a solid to a liquid.

  • Latent Heat of Vaporization: Heat required to change a liquid to a gas.

Density

• Density is the measure of how compact the mass in an object/substance is.

• Mass per unit volume

  • SI unit: kg/m³

  • ρ = m/v

    • ρ: density

    • m: mass

    • v: volume

Buoyancy

• Buoyancy is the upward force exerted by a fluid (liquid or gas) on an object immersed in it.

• This force enables objects to float or appear lighter when submerged in a fluid.

• Buoyancy is a direct consequence of pressure differences within the fluid.

Archimedes’ Principle

• Discovered by the ancient Greek mathematician and engineer Archimedes.

• Archimedes’ Principle states: "Any object completely or partially submerged in a fluid experiences an upward buoyant force equal to the weight of the fluid displaced by the object."

Atomic and Molecular Interactions

• Intermolecular Forces: Forces between molecules.

  • Van der Waals forces: Weak attractions between all atoms and molecules.

  • Dipole-Dipole interactions: Attractions between polar molecules.

  • Hydrogen bonding: Strong type of dipole-dipole interaction involving hydrogen and electronegative atoms like O, N, or F.

• Chemical Bonds:

  • Ionic Bonds: Transfer of electrons from one atom to another.

  • Covalent Bonds: Sharing of electron pairs between atoms.

  • Metallic Bonds: Delocalized electrons shared among metal atoms.

Crystalline and Amorphous Solids

• Crystalline Solids: Atoms arranged in a highly ordered structure, forming a crystal lattice (e.g., salts, diamonds).

• Amorphous Solids: Lack a long-range order in atomic structure (e.g., glass, plastics).

Properties of Materials:

• Mechanical Properties: Strength, ductility, hardness, toughness.

• Thermal Properties: Conductivity, expansion.

• Electrical Properties: Conductivity, resistivity.

• Composite Materials: Combinations of two or more materials to enhance properties (e.g., fiberglass, concrete).