3.1.1.3 Electron configuration- First ionisation energy Flashcards
why the first ionisation energy of rubidium is less than the first ionisation energy of krypton
outer electron in Rb is in 5th shell (or additional shell) further away (or more shielded) from nucleus
Chemical equation for the enthalpy change: the first molar ionisation energy of silicon
Si(g) → Si+ (g) + e–
Chemical equation for the enthalpy change: the second molar ionisation energy of silicon
Si+ (g) → Si2+(g) + e–
reasons why the first ionisation energy of neon is lower than the third ionisation energy of magnesium
Mg2+ ion smaller than Ne atom / Mg2+ e – closer to nucleus
Mg2+ has more protons than Ne / higher nuclear charge
The general trend in the first ionisation energies of the Period 3 elements, Na – Ar
trend: increases
more protons / increased proton number / increased nuclear charge
same shell / same shielding / smaller size
why the first ionisation energy of sulphur is lower than would be predicted from the general trend of period 3 elements
e- pair in the 3p sub-level
repulsion between the e– in this e– pair
Definition: first ionisation energy
Energy to remove 1 electron
from a gaseous atom
why boron has a lower first ionisation energy than beryllium
Be’s outer electron is in an s (2s) orbital
B’s outer electron is in a p (2p) orbital
B’s outer electron is higher in energy
why the first ionisation energy of helium is very large
Electron is not shielded from nucleus
the general trend in the first ionisation energy of the Period 3 elements from Na to Ar
Increases
how, and explain why, the first ionisation energy of aluminium does not follow this general trend
lower than expected / lower than Mg / 1 less energy needed to ionise; e– removed from (3)p sub-level;
of higher energy / further away from nucleus / shielded by 3s e – s;
an equation, including state symbols, to represent the process for which the energy change is the first ionisation energy of Na
Na(g) → Na+ (g) + e–
the general trend in the values of the first ionisation energies of the elements Na to Ar
Trend : Increases Explanation : Increased nuclear charge or proton number Stronger attraction (between nucleus and (outer) e– )
how, and explain why, the values of the first ionisation energies of the elements Al and S deviate from the general trend
How values deviate from trend: (both values) too low
Explanation for Al: e– removed from 3p
e – or orbital is higher in energy level or better shielded than (3)s
or p electron is shielded by 3s electrons
Explanation for S: e– removed from (3)p electron pair (1) repulsion between paired e– (reduces energy required)
Definition: first ionisation energy
Heat / enthalpy / energy for removal of one electron (1) from a gaseous atom
Why the first ionisation energy of neon is higher than sodium
Neon’s electron is in a lower (2p) shell
attracted more strongly to (or less shielded from) the nucleus
Why the first ionisation energy of magnesium is higher than sodium
more protons
electrons in same shell
similar shielding
why the first ionisation energy of aluminium is lower than that of magnesium
Als outer electron is in a 3p sub-shell
higher in energy than 3s in Magnesium
the variation in first ionisation energy of the elements across Period 3 from sodium to argon
general increase across period
because number of protons (or nuclear charge) increases
but electrons in same shell (or similar shielding)
fall from Mg to Al
Al’s outer electron is in a p orbital
higher in energy than s electron in Mg
fall from P to S
two of the p electrons in S are paired (or in same orbital)
the trend in the first ionisation energies of the elements in Group 2 from magnesium to barium
Decrease
Ions get bigger / more (energy) shells
Weaker attraction of ion to lost electron
the element in Period 3 that has the highest first ionisation energy
Argon / Ar
Large(st) number of protons / large(st) nuclear charge
Same amount of shielding / same number of shells / same number of energy levels