F321 - Topic 1 Atoms and Electron Structure

  • Created by: Sophie
  • Created on: 26-01-14 14:18

Atomic Structure

Atomic Structure

  • The nucleus contains almost all of the mass of an atom, as it cotains protons and neutrons.
  • The nucleus of an atom contains all of the positive charge.
  • The electrons orbit outside the nucleus.
  • Electrons are negatively charged.


  • Atomic number = number of protons in the nucleus.
  • Mass number = number of protons and neutrons in the nucleus.
  • Number of neutrons = mass number - atomic number.
  • Number of electrons = number of protons (in a neutral atom)
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Isotopes and Ions


  • Isotopes are atoms with the same number of protons but different number of neutrons.
  • Isotopes of an element have the same chemical properties, as they have the same electron arangement.


  • Ions are formed when atoms gain or lose elctrons.
  • E.g. An atom Cl has 17 electrons, a Cl- ion has gained one electron.
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Relative Atomic Mass and Relative Isotopic Mass

Relative Atomic Mass

  • Average mass of an atom relative to 1/12 of the mass of a carbon-12 atom.
  • To calculate relative atomic mass, add together (mass number x percentage/100).


  • 75% of Cl atoms have a mass number of 35
  • 25% of Cl atoms have a mass number of 37
  • Average mass of a Cl atom    = (mass no x percent/100) + (mass no x percent/100)

                                                      = (35 x 75/100) + (37 x 25/100)

                                                      = 35.5

Relative Isotopic Mass

  • Mass of a particular isotope relative to 1/12 of the mass of a carbon-12 atom.
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  • An orbital is a region that can hold up to two electrons with opposite spins.
  • Orbitals have different shapes called - s, p, d and f.

S orbitals

  •  S orbitals are spherical in shape and come in sets of one.

P orbitals

  • P orbitals are hour - glass shaped and come in stes of three.

D orbitals

  • D orbitals come in sets of five which can hold up to 10 electrons.
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Energy Levels

E.g.    Iron - 26 electrons - 1s2 2s2 2p6 3s2 3p6 4s2 3d6

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Successive Ionisation Energies

Successive Ionisation Energies

  • Energy change when one mole of electrons is removed from one mole of gaseous positively charged ions.
  • In an element, successive ionisation energies get bigger because the emaining electrons are held more tightly by the unchanged nuclear charge.

First Ionisation Energy

  • Energy change when one mole of electrons is removed from one mole of gaseous atoms.
  • M(g) -----> M+(g) + e-
  • Ionisation is easier if :
    • The nuclear charge is smaller.
    • The electron is further away from the nucleus.
    • There is more shielding.
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First Ionisation Energies - First 20 Elements


  • First ionisation energy shows periodicity.

Factors affecting the size of ionisation energy

  • The size of that attraction will be governed by:
    • The charge on the nucleus.
    • The distance of the electron from the nucleus.
    • The number of electrons between the outer electrons and the nucleus (shielding).
    • Whether the electron is on its own in an orbital ( paired electrons are easier to remove).
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The Pattern in Periods 2 and 3


  • Why the ionisation energy increases:
    • There is an increasing number of protons in the nuceus.
    • The increasing nuclear charge, drags the outer electrons in closer to the nucleus.
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Trends in ionisation energy down a group


  • Why:
    • There is a bigger number of protons.
    • But, the shielding and nuclear radius overcomes the nuclear charge.
  • E.g.
    • Li -1s22s1(http://www.chemguide.co.uk/atoms/properties/padding.GIF)1st I.E. = 519 kJ mol-1
    • Na -1s22s22p63s1(http://www.chemguide.co.uk/atoms/properties/padding.GIF)1st I.E. = 494 kJ mol-1
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Blocks in the Periodic Table


  • The s block is group 1 and 2.
  • The p block is group 3 and 8.
  • The d block is between the s and p blocks.
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