The Atomic Structure

Overview of Atomic Structure

• Matter is composed of particles, which can be atoms, molecules, or ions.
• Atoms are extremely small and follow the Law of Conservation of Mass, which states that mass cannot be created or destroyed in a chemical reaction.

Subatomic Particles

• Atoms are made of smaller particles: protons, neutrons, & electrons
• These 3 particles are fundamental
• Atoms of different elements are made of different combinations of the three (that's why different elements have different relative masses!)

This table tells you everything you need to know about protons, neutrons, & electrons

• As you can see, the relative mass of an electron is negligible (Very small).
• This means most of the mass of an atom is concentrated in the nucleus.
• You find this right in the middle; the nucleus is only a tiny fraction of the atom's total volume.
• The incredibly high density of the nucleus suggests that the particles inside it are held very close together by an extremely powerful force.
• Protons are positively charged, so without this force the positive.

Historical Development of Atomic Theory

1. Dalton: Proposed the first atomic theory, suggesting that atoms are indivisible particles that combine in fixed ratios to form compounds.
2. Crookes: Investigated the properties of cathode rays using vacuum tubes, laying the groundwork for the discovery of the electron.
3. Stoney: Named the electron as the fundamental unit of electric charge.
4. Thomson: Discovered that cathode rays were composed of negatively charged particles, which he identified as electrons. He also measured the charge-to-mass ratio (e/m) of the electron.
5. Millikan: Measured the magnitude of the charge on the electron through the oil drop experiment.
6. Rutherford: Conducted the alpha-particle scattering experiment, leading to the discovery of the nucleus and proving that most of the atom is empty space. He also discovered protons in the nucleus.
7. Bohr: Proposed a model where electrons move in fixed orbits around the nucleus with specific energy levels.
8. Chadwick: Discovered the neutron, completing the model of the atom.

Isotopes and Atomic Mass

Isotopes are atoms of the same element with the same atomic number but different mass numbers due to varying numbers of neutrons.

• Example: Carbon-12 and Carbon-14 are both isotopes of carbon, differing in neutron number. Relative Atomic Mass ($A_r$):

• Defined as the weighted average of the mass numbers of all naturally occurring isotopes of an element, based on the carbon-12 scale.

• Example: The relative atomic mass of chlorine is approximately 35.5 due to the presence of isotopes chlorine-35 and chlorine-37 in different natural abundances.

Mass Spectrometry

Mass spectrometry is a technique used to determine the relative atomic mass of elements by measuring the mass of their ions. The process involves several key stages:

1. Vaporisation: The sample is vaporised.
2. Ionisation: Atoms in the sample are converted into positive ions.
3. Acceleration: These ions are accelerated in an electric field.
4. Separation: The ions are separated based on their mass-to-charge ratio in a magnetic field.
5. Detection: The ions are detected, and their abundance is measured.

The principle behind mass spectrometry is that charged particles moving in a magnetic field are deflected to different extents depending on their mass. Heavier ions are deflected less than lighter ones, allowing them to be separated and identified.