IB Chemistry HL

    Cards (100)

    • Avogadro's number
      6.02 * 10^23
    • Molar mass (M) units

      g / mol
    • Kelvin unit of temperature
      Celcius + 273

      STP in data booklet is 0 degrees
    • Combined gas law, for changes to environment of a fixed mass of gas
    • Ideal gas equation, for current conditions of a gas
      PV=nRT

      T in K
    • Equal volumes of different gases at the same temperature and pressure...
      Have equal numbers of particles
    • Solution
      homogeneous mixture of a liquid (solvent) and another substance (solute)
    • [Solute]
      concentration of the solute, mol/dm3 or mol/L
    • Decimeter (unit)
      mm < cm < dm < m
    • Isotope
      Atom with different number of neutrons, differs in physical properties that depend on mass (ie. density)
    • Orbital
      Region of space where 2 electrons (with opposite spin) can be found

      Probable electron density
    • S orbital
      sphere shape
      1 per energy level
    • P orbital
      dumbbell shape
      3 per energy level
    • What order to fill orbitals in when drawing orbital diagrams
      Aufbau principle — the triangle with diagonal lines, orbitals with lower energy are filled before those with higher energy

      Hund's rule — Every orbital in a sub-level is singly occupied with electrons of the same spin before any one orbital is doubly occupied
    • D orbital
      5 per energy level
    • F orbital
      7 per energy level
    • Electron configuration notation: 1s^2
      1 is energy level, s is orbital type, 2 is how many es in the sub-level
    • Remove electrons (for cation) from which orbital in electron configuration?
      Highest level/shell number! Not Aufbau-diagonals order
    • First ionization energy definition
      Energy required to remove one mole of electrons from one mole of gaseous atoms

      Energy required to transition electron from first energy level/shell to ∞
    • Ionization energy trend
      Increases to the top (less electron shielding) and right (increase in effective nuclear charge)
    • Successive ionization energies (first, second, etc)
      Should increase progressively (since larger proton:electron ratio as you go)

      A huge increase means you're trying to remove an electron from a noble gas electron configuration (aka a new energy level, not just new sub-level)
    • Elements in the same period have...
      outer electrons in the same energy level
    • Transition metals (definition, the excluded elements)
      must have an incomplete d sub-level in one or more of is oxidation states (aka not Zinc which is full, and not Scandium which is empty)
    • Common characteristics of transition metals
      Variable oxidation number (ion charge), coloured compounds, magnetic properties, form complex ions (a metal ion bonded to a ligand via a covalent bond)
    • Effective nuclear charge on an electron
      the nuclear pull experienced by the outer electrons in an atom, taking into account the shielding effect of inner full shells of electrons
    • Atomic radius trends
      increases to the bottom (just bigger nucleus and more electron shells) left (less effective nuclear charge on valence electrons, they're further away)
    • Is anion/cation bigger or smaller than normal atom (in terms of atomic radius)
      anion > atom since additional electron repulsion increases radius of outer shell

      cation < atom since you lost the outer shell completely
    • Isoelectronic
      Same electron configuration (ie. Na+ and Mg 2+)
    • Electron affinity (definition and trend)
      the energy released when one mole of electrons is added to one mole of gaseous atoms

      opposite trend to atomic radius (increases to the top right)
    • Electronegativity (definition and trend)
      a measure of the attraction of a nucleus for bonding electrons

      same trend as electron affinity, opposite trend as atomic radius — excluding noble gases
    • Melting point trends in metals
      decrease down group 1, increase in ionic radii reduces the force of the attraction between the M+ ions and the delocalized electrons
    • Melting point trends in non-metals
      increase down group 17, increase in the strength of London dispersion forces (because valence electrons held less tightly and form temporary dipoles more easily) with increasing number of electrons
    • Ligand
      An ion or molecule with an electron pair to donate to form a covalent (coordinate) bond

      Lewis base
    • Transition metal coloured complexes (why, when)
      When ligands bond to the central metal ion, repulsion between the electrons in the ligands and those in the d orbitals of the transition metal ion causes the five d orbitals to split into two sets of different energy (non-degenerate orbitals)

      electrons can transition from the lower set to the higher set of d orbitals by absorbing energy, which are known as d-d transitions — we see the complementary colour of the light wavelength absorbed
    • When do transition metals (coloured complexes) absorb shorter wavelengths of light
      ligands higher in the spectrochemical series produce a larger splitting of the d orbitals = takes more energy for electrons to make d-d transitions = absorbs shorter wavelengths of light
    • Ionic compounds (definition, melting pt, solubility)
      ions held together in crystal lattice structures (surrounded by ions of opposite charge) by ionic bonds, electrostatic attraction (mutual attraction between opposite charges)

      formula expressed by simplest ratio

      high melting point, soluble in polar solvents (water)
    • Are ionic compounds conductive
      conduct electricity when molten or aqueous but NOT when solid
    • What differences in electronegativity define a bond as non-polar covalent, polar covalent, or ionic?
      non-polar covalent = EN difference under 0.5
      polar covalent = EN difference between 0.5 and 1.7
      ionic = EN difference greater than 1.7
    • Covalent compounds (definition, what impacts bond length, what causes polarity, solubility)
      form by the sharing of 2 electrons

      Increasing number of bonds = shorter and stronger bonds

      Polar bonds form when the two atoms bonded together have different electronegativity values (more than 0.5 difference) (exception: if molecular shape results in cancellation of polar bonds)

      Like dissolves like — polar covalent soluble in polar solvent, non-polar covalent soluble in non-polar solvent
    • Coordinate covalent bond
      when both shared electrons are from the same atom
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