Ionisation energy - the amount of energy, required to remove one electron from one mole of a gaseous atom.
Electronegativity - Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons (measured on the Pauling Scale).
Atomic radius - The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the center of the nucleus to to the boundary of the surrounding cloud of electrons.
Nuclear charge - Nuclear charge is to do with the number of protons in the nucleus, i.e. the greater the number of protons the higher the nuclear charge.
Nuclear attraction - The attraction between the negative electrons in their orbitals, and the positive protons in the nucleus. The stronger the attraction the closer the electrons and protons become.
Shielding effect - Electrons in an atom can shield each other from the pull of the nucleus. This effect, called the shielding effect, describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. The more electron shells there are, the greater the shielding effect experienced by the outermost electrons.
The more shells there are, the moos shielding, therefore the first I.E decreases.
Trends or Periodicity
The periodicity of these properties follows trends as you move across a row or period of the periodic table or down a column or group:
Across a period:
Ionization Energy Increases
Atomic Radius Decreases
Down a group:
Ionization Energy Decreases
Atomic Radius Increases
“Explain why ionisation energy increases across a period, using period 2 and/or 3 as an example”
The charge on the nucleus - increases
The distance of the electron from the nucleus/ Atomic radius - decreases
The number of electrons between the outer electrons and the nucleus - increases
In the whole of period 2, the outer electrons…