as chem

• Subatomic Particles


  • Protons (P)
  • Neutrons (n)
  • Electrons (e-)
• Protons


  Relative Charge: +1, Relative mass/a.m.u: 1
• Neutrons


  Relative Charge: 0, Relative mass/a.m.u: 1
• Electrons


  Relative Charge: -1, Relative mass/a.m.u: 1/1840
• Mass concentrated within the centre, nucleus
• An atom is electrically neutral; P+ = e-
• Atomic no. or proton no. (Z)

  No. of protons
• Atomic mass or nucleon no. (A)

  No. of P + N
• Isoelectronic ions


  Ions having the same no. of e-s

  
• Isotopes


  Atoms of the same element with the same proton number but different numbers of neutrons
• Isotopes have similar chemical properties since they have the same number of protons and electrons (so chemical interactions are similar)
• Isotopes have different physical properties since they have different numbers of neutrons, causing them to have different masses and, therefore, different physical interactions
• Behaviour of a Beam of Subatomic Particles

  1. Protons: deflected to -ve pole

  2. Neutrons: not deflected

  3. Electrons: deflected to +ve pole

  4. Electrons: deflected at greater angle than Protons

  
• Principle quantum no. (P.Q)


  Describes each shell
• Subshells


  s, p, d, f
• Orbital


  Region in space where there is a maximum probability of finding an electron
• Each orbital can hold 2e-s in opposite directions
• When e-s are placed in a set of orbital of equal energy, they occupy them singly, and then pairing takes place
• 1. s placed in the opposite direction: both -ve charge & if placed in the same direction, they'd repel. In the opposite direction, they create a spin to reduce repulsion
• Completely filled or half-filled (i.e. one e- in each orbital) are more stable (reduced repulsion)
• In certain cases (e.g., period 3 elements), an electron would prefer the 4s orbital over 3d while filling up
• When losing electrons, the electrons from the 4s orbital would be lost first, and then those from the 3d orbital would be lost
• Aufbau's Principle


  Method of showing how atomic orbitals are filled in a definite order to give the lowest energy arrangement possible
• Energy difference between s & p orbitals is very small - an e- from p can be promoted to half-fill or full-fill p orbital to make the atom more stable
• s orbitals


  Spherical, with the nucleus at the centre
• Free radical


  Species with one or more unpaired electrons
• Ionisation Energies (I.E)


  1st I.E: the energy needed to remove 1 mole of e-s from 1 mole of a gaseous atom to form 1 mole of unipositive ions
• Each Successive I.E is higher than the previous one because as e-s are removed, protons > e-s - the attraction between protons and remaining electrons increases
• Successive I.Es have a large jump in their value when e-s removed from the lower energy shell
• Deducing group no.


  By checking when 1st big jump in I.E occurs
• Factors Affecting Ionisation Energy

  • Nuclear Charge: Greater nuclear charge means greater ionization energy
  • Shielding Effect: Greater effect lower I.E because lesser attractive force between nucleus & outer e-s
  • Atomic Radius: Greater radius lower I.E; a distance of outermost e- to the nucleus is large - less energy needed to remove e-
  • Stable Configuration: High I.E needed to remove e-s from completely or half-filled orbitals
• I.E of Al lower than Mg: e- removed in Al is from higher energy 3p orbital which is further away from nucleus than 3s e- being removed from Mg. Nuclear attraction is less for 3p than 3s - I.E of Al is lower than Mg
• I.E of S lower than P: e- being removed in P is in a half filled, more stable 3p orbital whereas in S, the pairing of electrons in 3p results in increased repulsion - less energy need to remove an e-

  
• Ionic Radius


  Describes the size of an ion
• Ionic Radius Trends


  • Positive Ion: smaller radius than original neutral atom because shell no. decreases, the screening effect decreases, but the attraction of the nucleus increases
  • Negative Ion: larger ionic radius than neutral atom because e-s added while nuclear charge remains same
  • Groups 1 to 3: Positive Ions, Groups 5 to 7: Negative Ions
  • Across the period: Ionic radius decreases as proton no. and effective nuclear charge increases
  • Negative ions are always larger than positive ions in the same period as they have one more shell
  • Ionic radius increases down the group since the number of electron shells increases
  • As the negative charge on anion increases, the ionic radius increases since the number of electrons gained increases such that the number of electrons exceeds the number of protons
  • As the positive charge on the cation increases, the number of electrons lost increases, so the electrostatic attraction between the nucleus and outer electrons increases
• Relative Mass


  • Atomic mass (Ar): weighted average mass of an atom
  • Molecular mass (Mr): mass of a molecule
  • Formula mass: mass of one formula unit of a compound
  • Isotopic mass: mass of a particular isotope of an element
• Unified atomic mass unit

  u = 1.66 x 10-27kg
• Mole


  Amount of substance that has the same number of particles (atoms, ions, molecules or electrons) as there are atoms in exactly 12g of the carbon-12 isotope
• Avogadro's constant


  Number of atoms, ions, molecules or electrons in a mole = 6.02 x 10^23
• Mass Spectra


  Abundance of isotopes can be represented on a mass spectra diagram