Particles Revision Notes

Understanding Atomic and Molecular Scale

Atoms and small molecules are microscopic particles, meaning they cannot be observed directly with the human eye. Their dimensions are so small that special units of measurement and scientific notation are used to describe them. This scale is fundamental to understanding the properties and interactions of matter at a fundamental level. It's important to grasp the relative sizes rather than memorising exact values for every single atom or molecule.

Typical Size of an Atom

The typical diameter of an atom is approximately 0.1 nanometres (nm). To express this in metres, which is the standard SI unit of length, we use the conversion factor 1 nm = 1 x 10⁻⁹ m. Therefore, an atom's diameter is about 0.1 x 10⁻⁹ m, which can also be written as 1 x 10⁻¹⁰ m. This extremely small size means that even a tiny speck of dust contains billions of atoms.

Typical Size of Small Molecules

Small molecules are formed when a few atoms bond together. Consequently, their size is generally slightly larger than a single atom but still within the same order of magnitude. Small molecules typically have a diameter ranging from approximately 0.1 nm to 1 nm. Examples include water (H₂O) or carbon dioxide (CO₂). This range is still incredibly small and necessitates advanced techniques for observation.

Units of Measurement and Powers of Ten

When discussing the size of atoms and molecules, nanometres (nm) are the most convenient unit. Remember that 'nano' means one billionth (10⁻⁹). So, 1 nanometre = 10⁻⁹ metres. This helps in understanding scientific notation like 1 x 10⁻¹⁰ m for an atom's diameter, which is equivalent to 0.1 nm. Being comfortable with these powers of ten is crucial for describing microscopic dimensions.

Dalton's Indivisible Spheres (Early 19th Century)

John Dalton proposed that atoms are tiny, indivisible, solid spheres. He believed that each element consisted of identical atoms with the same mass, and atoms of different elements had different masses. This model helped explain the conservation of mass and the ratios in which elements combine, but it did not account for any internal structure within the atom. It was the first widely accepted atomic theory that laid the groundwork for future discoveries.

Thomson's Plum Pudding Model (Late 19th Century)

J.J. Thomson discovered the electron through cathode ray experiments. This led him to propose the 'plum pudding' model, where the atom was described as a ball of positive charge with negatively charged electrons (the 'plums') embedded within it. This model was a significant step as it was the first to suggest that atoms were not indivisible but contained smaller, charged subatomic particles.

Rutherford's Nuclear Model (Early 20th Century)

Ernest Rutherford's alpha particle scattering experiment provided crucial evidence against the plum pudding model. Most alpha particles passed directly through a thin gold foil, but a small proportion were deflected at large angles, and some even bounced back. This scattering pattern suggested that the atom's mass and positive charge were concentrated in a tiny central nucleus, with electrons orbiting around it. This nuclear model fundamentally changed the understanding of atomic structure.

Bohr's Energy Levels Model (Early 20th Century)

Niels Bohr refined Rutherford's model by proposing that electrons orbit the nucleus in specific, fixed energy shells or energy levels. He suggested that electrons could move between these shells by absorbing or emitting specific amounts of energy. This model successfully explained the emission spectra of hydrogen and why electrons do not spiral into the nucleus. It marked a transition towards modern quantum theory, providing a more stable and predictable atomic structure.

States of Matter: Solids

In a solid, particles are arranged in a regular, closely packed pattern. They are held together by strong forces of attraction. The particles vibrate about fixed positions but cannot move from one place to another. This arrangement gives solids a definite shape and volume, making them difficult to compress. Examples include ice and iron.

States of Matter: Liquids

In a liquid, particles are still closely packed but are randomly arranged. The forces of attraction between particles are weaker than in solids, allowing them to move past each other. Liquids have a definite volume but no fixed shape, as they flow to fill the bottom of a container. They are also difficult to compress. Examples include water and oil.

States of Matter: Gases

In a gas, particles are widely spaced and move randomly and rapidly in all directions. The forces of attraction between particles are very weak, almost negligible. Gases have no definite shape or volume and will spread out to fill their container. They are easy to compress. Examples include steam and air.

Changes of State

Changes of state occur when a substance is heated or cooled, leading to changes in particle arrangement and movement. Melting is the change from solid to liquid, where particles gain enough energy to overcome some forces and move freely. Boiling (or evaporation) is the change from liquid to gas, where particles gain enough energy to completely overcome forces and escape. Freezing is the change from liquid to solid, where particles lose energy and settle into fixed positions. Condensation is the change from gas to liquid, where particles lose energy and move closer together.

Atomic Number (Z) and Protons

The atomic number (Z) is a fundamental identifier for an element and is always equal to the number of protons found in the nucleus of an atom. This number dictates the element's chemical properties and its position on the Periodic Table. All atoms of a specific element have the same number of protons. For example, all carbon atoms have 6 protons, so their atomic number is 6. The atomic number is usually found as the smaller number associated with an element symbol, often written as a subscript preceding the symbol (e.g., ₆C).

Mass Number (A) and Neutrons

The mass number (A) represents the total count of protons and neutrons within an atom's nucleus. It is always a whole number. To calculate the number of neutrons in a specific atom, subtract the atomic number (number of protons) from the mass number: Number of neutrons = Mass number - Atomic number. Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons, leading to different mass numbers. For instance, Carbon-12 has 6 protons and 6 neutrons (mass number 12), while Carbon-14 has 6 protons and 8 neutrons (mass number 14). The mass number is often written as a superscript preceding the element symbol (e.g., ¹²C or ¹⁴C).

Electrons in Neutral Atoms and Ions

In a neutral atom, the number of electrons is exactly equal to the number of protons. This balance ensures that the atom has no overall electrical charge. When an atom gains or loses electrons, it becomes an ion, carrying a net electrical charge. Losing electrons forms a positive ion (cation); the number of electrons decreases. Gaining electrons forms a negative ion (anion); the number of electrons increases. The number of protons and neutrons remains constant when an atom forms an ion. For example, an oxygen atom (8 protons) has 8 electrons. An O²⁻ ion (still 8 protons) has gained two electrons, so it has 10 electrons.