Chem Mod 1 Ch 3 Atomic Structure and Mass

• Electrostatic attraction


  The attraction between negatively-charged and positively-charged particles
• Isotope


  A different form of that element, with the same number of protons but a different number of neutrons
• Radioisotope


  An unstable form of a chemical element that releases radiation as its nucleus breaks down to become stable
• Atomic models


  • Atoms initially thought of as hard spheres but are mostly empty space
  • Made up of smaller/subatomic particles: protons and neutrons in positively charged nucleus
  • Electrons form a cloud around the nucleus and orbit in specific parts of the cloud which is negatively charged (electron mass is negligible)
• Subatomic particles


  • Proton
  • Neutron
  • Electron
• Scientific notation


  a × 10^b
• Atomic structure


  • A = mass number (# of protons + neutrons)
  • Z = atomic number (# of protons)
  • X = symbol of element
• Isotopes


  Atoms that have the same number of protons (atomic number) but different numbers of neutrons (and therefore different mass numbers)
• Radioisotopes


  Isotopes that are radioactive
• Stable and unstable isotopes


  • Stable isotopes lie within the band of stability where the nuclei have the right balance of protons and neutrons
  • Unstable isotopes have either too many or too few neutrons to be stable, meaning they are radioactive
• Radiation, decay and nuclear equations


  1. Radioisotopes break down (decay) to form stable nuclei, during which, radiation is emitted in the form of alpha, beta or gamma particles

  2. Alpha particles are the least penetrable whilst gamma are the most

  3. Alpha particles are emitted by nuclei that have too few neutrons to be stable

  4. Beta particles are emitted by nuclei that have too many neutrons to be stable

  5. Gamma radiation is a type of high-energy electromagnetic radiation

  6. When balancing nuclear equations, the atomic and mass numbers must add up equally on both sides
• Mass spectrometer


  • Measures relative isotopic masses of elements and their isotopic abundances
  • Ionisation: The atom/molecule is ionised by knocking one or more electrons off to make a positive ion
  • Acceleration: The ions are accelerated so that they all have the same kinetic energy
  • Deflection: The ions are then deflected by a magnetic field according to their masses
  • Detection: The beam of ions passing through the machine is detected electrically
• Mass spectrum


  • The output from the mass spectrometer
  • Number of peaks = number of isotopes
  • Horizontal (x) axis = relative mass of each isotope present in the sample
  • Vertical (y) axis = % relative abundance of each isotope in the sample
• Relative atomic mass (Ar)


  The average value of the isotopes of the element, taking into account the percentage abundance of each isotope
• Relative isotopic abundance / percentage abundance


  % abundance = peak height / total peak height × 100
• Relative molecular mass


  The sum of the relative atomic masses of each atom in the molecule
• Relative formula mass


  The sum of the relative atomic masses of the elements in the formula (for compounds that do not exist as molecules)
• Flame tests


  • Used to determine the identity of a metal in a sample by heating it, resulting in the metal giving off a characteristic colour
• Emission spectra


  • When atoms are heated, they give off electromagnetic radiation/light, which, when passed through a prism produces a line spectra/emission spectrum
  • An emission spectrum is a black background with coloured lines which corresponds to a different type of electromagnetic radiation
  • As each element has a unique emission spectrum, it can be used to identify an element
• Bohr model of the atom


  • Electrons revolve around the nucleus in fixed, circular orbits
  • The electrons' orbits correspond to specific energy levels in the atom
  • Electron can only occupy fixed energy levels and cannot exist between two energy levels
  • Orbits of larger radii correspond to energy levels of higher energy
• Electron shells


  Electrons are grouped in different energy levels, called electron shells, each labelled with the number n = 1, 2, 3,…, with n = 1 being the closest to the nucleus
• Emission spectra and the shell model


  1. Heating an element can cause an electron to absorb energy and jump to a higher energy level (excited state)

  2. Shortly after, as its energy decreases, the electron returns to the lower energy level, releasing a fixed amount of energy as light (the difference between the 2 energy levels)

  3. Each transition corresponds to a specific energy of light and thus, one line in the line spectrum for that element
• Ionisation energy


  The energy needed to remove an electron from an atom (the first electron to be removed has the lowest ionisation energy and thus, is the easiest to remove)
• Electronic configuration


  • The arrangement of electrons around the nucleus
  • Electrons are always as close to the nucleus as possible, which is why electrons typically occupy inner shells before outer ones
  • Each shell can contain a maximum number of electrons (shell n can hold 2n^2 electrons)
  • Lower energy shells fill before higher energy shells
  • Electron shells fill in a particular order
• Valence electrons

  • The electrons found in the valence shell (outermost shell)
  • Require the least energy to remove
  • Commonly involved in chemical reactions
  • Indicates chemical properties
• Pauli's exclusion principle


  • Each orbital can contain a max of 2 electrons, with each electron having a different spin
  • The total number of orbitals in a shell is given by n^2, resulting in a total of 2n^2 electrons per shell
  • An s-subshell has 1 orbital and can hold up to 2 electrons
  • A p-subshell has 3 orbitals and can hold up to 6 electrons
  • A d-subshell has 5 orbitals and can hold up to 10 electrons
  • A f-subshell has 7 orbitals and can hold up to 14 electrons
• Electronic configuration and the Schrödinger model


  • Coefficient represents the shell number (n)
  • Letters specify the subshell being filled
  • Superscripts denote how many electrons are in the subshell
  • Order of the subshells – 1s < 2s < 2p < 3s < 3p < 4s < 3d…
• Aufbau principle


  The lowest energy orbitals are always filled with electrons first
• Hund's rule


  Every orbital in a subshell must first be filled with one electron with the same spin before an orbital is filled with a second electron of opposite spin