Atomic structure

• Atoms are the basic building blocks of matter
• Atoms are composed of subatomic particles: protons, neutrons, and electrons
• Protons have a positive charge and are found in the nucleus of an atom
• Neutrons have no charge and are found in the nucleus of an atom
• Electrons have a negative charge and orbit the nucleus in electron shells
• The number of protons in an atom determines its atomic number
• The sum of protons and neutrons in an atom gives its atomic mass number
• Isotopes are atoms of the same element with different numbers of neutrons
• Electron configuration refers to the arrangement of electrons in an atom's electron shells
• The periodic table organizes elements based on their atomic number and chemical properties
• Chemical reactions involve the rearrangement of atoms to form new substances
• The relative isotopic mass is the mass of an atom of an isotope relative to 1/2 the mass of an atom of carbon 12
• The relative atomic mass is the weighted average mass of an atom of an element relative to 1/12 the mass of an atom of carbon 12
• the relative formula mass is the weighted average mass of a formula unit of a substance relative to 1/2 the mass of an atom of carbon 12
• the relative molecular mass is the weighted average mass of a molecule of a substance relative to 1/12 the mass of an atom of carbon 12.
• mass spectrometry is an analytical tool used to measure the mass to charge ratio (m/z) of one or more molecules present in a sample
• in mass spectrometry, the sample is ionised by bombarding it with electrons which removes an electron from some of the atoms in the sample creating positive ions.
• Electronic configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals
• Electrons fill orbitals starting from the lowest energy level to the highest
• Each orbital can hold a maximum of 2 electrons with opposite spins
• The electronic configuration of an atom is written using the Aufbau principle, Hund's rule, and the Pauli exclusion principle
• The Aufbau principle states that electrons fill the lowest energy levels first before moving to higher energy levels
• Hund's rule states that electrons will fill orbitals of the same energy level with parallel spins before pairing up
• The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers
• The electronic configuration of an atom can be represented using the electron configuration notation, which lists the energy levels and the number of electrons in each level
• For example, the electronic configuration of carbon (C) is 1s2 2s2 2p2, indicating the distribution of its 6 electrons in the 1s, 2s, and 2p orbitals
• The Heisenbergs uncertainty principle states that you cannot determine the position and momentum of an electron at the same time
• • S - spherical
  • p - dumb-bell shaped
  • d - various (donut)
  • f - various
• principal energy levels get closer together as you get further from the nucleus
• chromium will, to achieve a more stable arrangement of lower energy, promote one of the 4s electrons into the 3D shell

  6 unpaired electrons with lower repulsion
• copper to achieve a more stable arrangement of lower energy, will promote one of the 4s electrons into the 3D.
• the noble gases are unreactive because they have full outer shells
• when atoms gain or lose electrons they become charged particles called ions
• An orbital is a region of space where one is likely to find an electron
• Electron configuration is the distribution of electrons around the atom's nucleus.
• A subshell can hold up to 2n^2 electrons
• Ionisation energy is the energy required to remove an electron from a gaseous atom or ion
• Ionisation energy generally increases across a period and decreases down a group in the periodic table
• Factors affecting ionisation energy:

  • Nuclear charge: the greater the nuclear charge, the higher the ionisation energy
  • Distance from the nucleus: the closer the electron is to the nucleus, the higher the ionisation energy
  • Shielding effect: inner electrons shield outer electrons from the full nuclear charge, reducing ionisation energy
• First ionisation energy is the energy required to remove the outermost electron from a neutral gaseous atom to form a positive ion