Atomic Structure

• Relative isotopic mass - the mass of one atom of an isotope of an element relative to the mass of 1/12 of a carbon atom.
• Label

  A) Ionisation

  B) Acceleration

  C) Deflection

  
• Mass spectrometry steps 1-3

  1. Vapourise the sample of an element and inject it into the mass spectrometer.
  2. Ionisation - a high voltage is passed across the chamber, and due to electron bombardement, an electron is removed from atoms leaving 1+ charged ions.
  3. Acceleration - positive ions are now accelerated using an electric field generated by charged plates.
• Mass spectrometry steps 4-5

  4. Deflection - ions are deflected by an electromagnetic field into a curved path. Ions that are too heavy (high m/z value) won't be deflected enough, and ions that are too light are deflected too much.

  5. Charged Particle Detector - ions hit the detector, producing a smaller current as electrons move from the detector to the positive ion. The size of the current produced is proportional to the abundance of the species.
• An equation that represents what is occurring during the Ionisation step during mass spectrometry is:

  M(g) + e- = M+ + e- + e-

  (In the products, one of the electrons is from the voltage and one from the atom)
• An equation that represents what occurs at the charged particle detector during mass spectrometry is:

  M+ + e- = M(g)
• Mass spectrometry takes place in a vacuum because:

  • it prevents ions from colliding with air particles.
  • it removes traces of other samples.
  • it prevents air particles being detected instead.
• Orbital - the volume of space around an atom that can contain a maximum of 2 electrons.
• shells are made up of sub-shells
• the s sub-shell contains 1 orbital and thus can hold a max of 2 electrons.

  the p sub-shell contains 3 orbitals and thus can hold a max of 6 electrons.

  the d sub-shell contains 5 orbitals and thus can hold a max of 10 electrons.

  the f sub-shell contains 7 orbitals and thus can hold a max of 14 electrons.
• Shells

  first shell only has an s sub-shell and thus can hold a max of 2 electrons.

  second shell has an s and p sub-shell and thus can hold a max of 8 electrons.

  third shell has an s, p, and d sub-shell and thus can hold a max of 18 electrons.

  fourth shell has an s,p,d, and f sub-shell, and thus can hold a max of 32 electrons.
• An S orbital looks like this

  A) .

  
• P orbitals look like these:

  A) P(x)

  B) P(y)

  C) P(z)

  
• Below is the order of energy for spdf notation:

  
• isotopes have the same chemical reactions because they have the same electronic configuration as eachother.
• Spin Diagrams - electrons can spin either up or down. An orbital is made up of pairs of the oppositely spinning electrons
• Rules for filling spin diagrams

  • electron pairs in the same box/orbital must have opposite spins.
  • If a sub-shell can hold more than 1 pair of electrons/orbital (p and d), electrons should be put in different boxes/orbitals pointing the same way, before pairing up in an orbital with an already present electron. In other words, electrons should be filled in empty degenerate (of equal energy) boxes before pairing up.
• Below is the spin diagram for oxygen:

  
• Atomic radius trends:

  Across a period - decreases

  Down a group - increases
• Atomic radius trend explanation (decreases across a period)

  • the nuclear charge increases.
  • electrons are entering the same shell, thus the number of shells remains constant and there's no increase in shielding of core electrons.
  • Theres a stronger electrostatic attraction between the nucleus and the valence electrons, which decreases the atomic radius.
• Atomic radius trend explanation (increases down a group)

  • the increase in nuclear charge is cancelled out by the increase in shielding done by the core electrons.
  • The atomic radius thus increases as valence electrons are in a higher energy shell.
  • this is because there's now a weaker electrostatic attraction between the nucleus and valence electrons, which increases the atomic radius.
• First ionisation energy - the enthalpy change when 1 mol of electrons is removed from 1 mol of atoms forming 1 mol of ions in the gas phase.
• First Ionisation Energy Trends:

  Across a period - increases (harder to remove electron = higher energy)

  Down a group - decreases (easier to remove electron = lower energy).
• First Ionisation Energy Trend Explanation (increases across a period)

  • nuclear charge increases.
  • shielding remains constant
  • atomic radius decreases due to increased electrostatic attraction of the nucleus to the valence electrons.
  • Thus harder for a valence electron to be removed.
• First Ionisation Energy Trend Explanation (decreases down the group)

  • the atomic radius increases.
  • thus outermost electrons are held successively further from the nucleus.
  • thus outer electrons become more shielded from attraction of the nucleus by repulsion of inner electrons, despite the increase in nuclear charge.
  • thus its easier for outer electrons to be lost.
• Succesive ionisation energies are always larger
• Successive ionisation energies explanation:

  • when the first electron is removed, a positive ion forms.
  • this increases the proton:electron ratio, as the same number of protons is attracting fewer electrons.
  • Thus the ion gets smaller, and electrons being removed are more attracted to the nucleus.
  • thus the energy required to remove the next electron is higher, as the attraction between the nucleus and remaining electrons increases.
• For Carbon, the 5th ionisation energy is MUCH greater than the 4th. This shows that carbon is in group 4, as there are 4 electrons in its outer shell, and the 5th one being removed is in a lower energy shell (closer to the nucleus) thus making it harder to be removed.
• Primary factors affecting ionisation energy:

  • nuclear charge
  • atomic radius
  • shielding
• Secondary factors affecting ionisation energy:

  • energy of the sub-shell
  • occupancy of the orbital
• IE trend exceptions:

  Across a period, there is normally an increase in IE.

  However from groups 2-3, there is a small decrease because:

  • you enter a new sub-shell
  • the valence electron in your group 3 atom is in a higher energy sub-shell. This is caused by the shielding of the sub-shell prior.
  • This factor outweighs the increase in nuclear charge, and its thus means the electrostatic attraction between the nucleus and valence electrons decreases.
• IE trend exceptions:

  Across a period, there is normally an increase in IE.

  However from groups 5 -6 there is a small decrease because:

  • one orbital is now doubly occupied.
  • this causes electron:electron repulsion.
  • this factor outweighs the increase in nuclear charge. Thus the electrostatic attraction between the nucleus and the valence electrons decreases.
• Electronic configuration exceptions:

  Cr + Cu make either a half or fully filled d subshell

  Cr: [Ar] 4s1 3d5

  Cu: [Ar] 4s1 3d10
• Electronic configuration exceptions:

  Transition metal ions (d block) are always 4s0

  Fe2+: [Ar] 4s0 3d6

  Ni2+:[Ar] 4s0 3d8

  Fe3+:[Ar] 4s0 3d5
• in mass spectrometry, the largest m/z peak is due to the complete molecule, and is equal to the Mr of the molecule. This peak is known as the molecular ion peak.

  
  

  A molecule fragmenting will cause it to produce a positively charged fragment, and a radical.

  
• in a mass spectrum, there may be a small M+1 peak present. This is caused by a single Carbon 13 isotope being present.

  
• relative atomic mass - weighted mean mass of an atom relative to the mass of 1/12th of an atom of Carbon 12
• relative formula mass - The sum of the relative atomic masses (Ar) of all atoms in a chemical formula.
• relative molecular mass - the weighted mean mass of a molecule relative to the mass of 1/12 of an atom of carbon-12. It is the sum of the relative atomic masses of each atom within the molecule.
• periodic table: