2 - Atomic Structure

• Protons and neutrons are located in the centre of the atom (the nucleus), while the electrons orbit the nucleus in shells.
• Protons have relative masses of 1, while the electron has a relative mass of 5 × 10^4.
• The mass number (A) is the number of protons and neutrons in the nucleus.
• The atomic number (Z) is the number of protons in an atom, which equals the number of electrons and is used to define which element the atom belongs to.
• Isotopes are atoms of the same element with the same atomic numbers but different mass numbers.
• Different isotopes have the same chemical properties, but different physical properties such as mass and density, fixed points, rates of diffusion.
• Radioisotopes are unstable and undergo radioactive decay.
• Cobalt-60 is used in radiotherapy, where the radiation destroys cells.
• Iodine-131 and iodine-125 are used as medical tracers, where after being injected into someone they release gamma rays which can be detected.
• Carbon-14 is used in radiocarbon dating, where the amount of 14 C in a sample is compared to the amount of 14 C in the air.
• The mass spectrometer works by vaporising the sample, ionising atoms, accelerating ions, deflecting ions, and detecting ions with a particular mass/charge (m/z) ratio.
• In a vacuum, the frequency of light is inversely proportional to the wavelength, meaning a continuous spectrum is made up of all the frequencies, while a line spectrum only contains certain distinct frequencies.
• Line spectra converge at high energy.
• Electrons can only exist in fixed energy levels.
• Supplying energy makes electrons go up a level, and they emit energy when they drop a level.
• The energy gap corresponds to a particular frequency of light, meaning drops to n = 1 are UV, to n = 2 are visible, to n = 3 are IR.
• Within a shell, there is a general increase because the effective nuclear charge increases (same number of protons but fewer electrons).
• The jumps between 1-2, 9-10 and 17-18 suggest the existence of main energy levels.
• The increase within the period is not linear, which may suggest sub-levels.
• Li, Na, and K have much lower ionisation energy than the elements before, showing that there are main shells.
• The general increase in ionisation energy is because the nuclear charge is increased while electrons are going into the main shell.
• B is lower than Be because the nuclear charge is partially shielded by the full 2s sub-shell.
• O is lower than N, because in N the 2p shell is half full.
• The extra electron in O is repelled by one of the electrons already in the 2p orbital, so it has lower ionisation energy.
• c = fλ.
• E = hf.
• The Aufbau principle states that electrons enter the lowest energy orbital available.
• The Pauli exclusion principle states that each orbital can hold a maximum of two electrons with different spins.
• Hund’s rule states that electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same level are available.
• Electron orbitals are often represented as a box, with arrows representing electrons.
• Longer electronic configurations can be expressed using noble gasses for abbreviations.
• The 4s and 3d subshells are very close to each other, and when a transition ion is formed, the 4s electrons are always lost first.
• When filling up shells, the 4s fills before the 3d.
• Copper has arrangement [Ar] 4s 1 3d 10 instead of [Ar] 4s 2 3d 9.
• Chromium has arrangement [Ar] 4s 1 3d 5 instead of [Ar] 4s 2 3d 4.
• This is because the d orbital is most stable when full, empty or half full.
• The first electron is easiest to remove because it is the furthest away from the nucleus and has the most shielding.