ionisation energy

Define first ionisation energy.

The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. Units: kJ mol⁻¹.

Define successive ionisation energy.

A measure of the energy required to remove each electron in turn from an atom.

Define second ionisation energy.

The energy required to remove one electron from each 1+ ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.

Write the equation for the first ionisation energy of sodium.

Na(g) → Na⁺(g) + e⁻.

Write the equation for the first ionisation energy of oxygen.

O(g) → O⁺(g) + e⁻.

Write the equation for the first ionisation energy of aluminium.

Al(g) → Al⁺(g) + e⁻.

Write the equation for the first ionisation energy of chlorine.

Cl(g) → Cl⁺(g) + e⁻.

Write the equation for the first ionisation energy of lithium.

Li(g) → Li⁺(g) + e⁻.

Write the equation for the first ionisation energy of magnesium.

Mg(g) → Mg⁺(g) + e⁻.

Write the equation for the first ionisation energy of bromine.

Br(g) → Br⁺(g) + e⁻.

Write the equation for the first ionisation energy of neon.

Ne(g) → Ne⁺(g) + e⁻.

Write the equation for the second ionisation energy of sodium.

Na⁺(g) → Na²⁺(g) + e⁻.

Write the equation for the second ionisation energy of calcium.

Ca⁺(g) → Ca²⁺(g) + e⁻.

Write the equation for the second ionisation energy of strontium.

Sr⁺(g) → Sr²⁺(g) + e⁻.

Write the equation for the third ionisation energy of aluminium.

Al²⁺(g) → Al³⁺(g) + e⁻.

Write the equation for the third ionisation energy of oxygen.

O²⁺(g) → O³⁺(g) + e⁻.

Write the equation for the fourth ionisation energy of magnesium.

Mg³⁺(g) → Mg⁴⁺(g) + e⁻.

Write the equation for the fourth ionisation energy of copper.

Cu³⁺(g) → Cu⁴⁺(g) + e⁻.

Write the equation for the seventh ionisation energy of sodium.

Na⁶⁺(g) → Na⁷⁺(g) + e⁻.

Write the equation for the seventh ionisation energy of bromine.

Br⁶⁺(g) → Br⁷⁺(g) + e⁻.

State the three factors that affect ionisation energy.

Nuclear charge (number of protons), distance of outermost electron from the nucleus (atomic radius), and electron shielding (from inner-shell electrons).

Explain how nuclear charge affects ionisation energy.

The more protons in the nucleus, the greater the nuclear charge. The greater the nuclear charge, the stronger the nuclear attraction on the outer electrons. Therefore a greater nuclear charge means more energy is needed to remove the outer electron.

Explain how distance (atomic radius) affects ionisation energy.

As the distance between the nucleus and the outermost electron increases, the attraction between them decreases. The weaker the nuclear attraction, the less energy is needed to remove the electron.

Explain how electron shielding affects ionisation energy.

Electron shielding is the repulsion between electrons in different inner shells. This shielding effect reduces the net nuclear attraction from the positive nucleus on the outer-shell electrons. The more inner shells there are, the greater the shielding effect and the weaker the nuclear attraction experienced by the outer electrons.

State the mnemonic for factors affecting ionisation energy.

NASA - Nuclear charge, Atomic radius (distance), Shielding, Attraction. OR NDSA - Nuclear charge, Distance, Shielding, Attraction.

State the trend in first ionisation energy down a group.

First ionisation energy decreases down a group.

Explain the trend in first ionisation energy down a group.

The outer electron is in a different shell further from the nucleus so the atomic radius increases. The outer electron experiences more shielding. The increased shielding outweighs the increased nuclear charge. The outer electron is less strongly attracted to the nucleus, so less energy is needed to remove the electron.

Explain the trend in first ionisation energy down a group and the effect on reactivity of metals.

First ionisation energy decreases down a group. The outer electrons are less strongly attracted to the nucleus and are more easily lost. Therefore reactivity of metals increases down a group.

Explain the trend in first ionisation energy down a group and the effect on reactivity of non-metals.

First ionisation energy decreases down a group. The outer electrons are less strongly attracted to the nucleus and are less easily gained. Therefore reactivity of non-metals decreases down a group.

State the general trend in first ionisation energy across a period.

First ionisation energy generally increases across a period.

Explain the general trend in first ionisation energy across a period.

The number of protons increases, so the nuclear charge increases. The outer electrons are in the same shell so experience the same shielding. The atomic radius decreases. The outer electron is more strongly attracted to the nucleus, so more energy is needed to remove the electron.

Explain why the first ionisation energy of beryllium is higher than lithium.

Both are in the same period so the outer electrons are in the same shell with the same shielding. Beryllium has one more proton than lithium and a smaller atomic radius. Beryllium has a stronger attraction of the nucleus to the outer shell electron. More energy is needed to remove the outer electron in beryllium.

Explain why the first ionisation energy of neon is higher than sodium.

Sodium has one more shell than neon. Sodium has a bigger atomic radius and more shielding from inner shell electrons. The increased distance from the nucleus and shielding effect far outweigh the increase in nuclear charge. There is a weaker attraction of the nucleus for the outer shell electron in sodium, so less energy is needed to remove the outer electron in sodium.

Explain why the first ionisation energy of sodium is lower than magnesium.

The outer electrons in sodium and magnesium are in the same shell and sub-shell so have the same shielding. Sodium has one less proton than magnesium and a bigger atomic radius. Sodium has a weaker attraction of the nucleus for the outer shell electron, so less energy is needed to remove the outer electron in sodium.

Explain why helium has the highest first ionisation energy.

Both helium and hydrogen have their outer electron closest to the nucleus in shell 1 with no shielding from inner shell electrons. Helium has one more proton than hydrogen. Helium has the strongest attraction of the nucleus for its outer electron. Helium requires the most energy to remove its outer electron.

Explain why francium is the most reactive Group 1 metal.

Francium has the most shells, the biggest atomic radius and the most shielding from inner shell electrons. The increased distance from the nucleus and shielding effect far outweigh the increase in nuclear charge. There is the weakest attraction of the nucleus to the outer electron in francium, so the least energy is needed to remove the outer electron.

Explain why the first ionisation energy of boron is less than beryllium.

In beryllium, the outer electron is removed from the 2s sub-shell. In boron, the outer electron is removed from the 2p sub-shell. The 2p sub-shell has a higher energy than the 2s sub-shell. Less energy is needed to remove the 2p electron from boron.

Explain why the first ionisation energy of aluminium is less than magnesium.

In magnesium, the outer electron is removed from the 3s sub-shell. In aluminium, the outer electron is removed from the 3p sub-shell. The 3p sub-shell has a higher energy than the 3s sub-shell. Less energy is needed to remove the 3p electron from aluminium.

Explain why the first ionisation energy of oxygen is less than nitrogen.

In nitrogen, each p orbital contains one unpaired electron. In oxygen, one p orbital contains two paired electrons. The paired electrons in oxygen repel each other. Less energy is needed to remove one of the paired electrons from oxygen than from nitrogen.

Explain why the first ionisation energy of sulfur is less than phosphorus.

In phosphorus, each p orbital contains one unpaired electron. In sulfur, one p orbital contains two paired electrons. The paired electrons in sulfur repel each other. Less energy is needed to remove one of the paired electrons from sulfur than from phosphorus.

Explain why successive ionisation energies always increase.

Each time an electron is removed, there are the same number of protons but fewer electrons. The proton:electron ratio increases. The remaining electrons are being removed from a positive ion. The remaining electrons are more strongly attracted to the nucleus. More energy is needed to remove each electron in turn.

Explain how successive ionisation energies provide evidence for electron shells.

There is a large increase (big jump) in ionisation energy when an electron is removed from a new shell which is closer to the nucleus. This new shell has less shielding from inner electrons. The large increase shows a different shell/energy level.

Explain how successive ionisation energies can be used to determine the group of an element.

The group number is the number of electrons in the outer shell. The largest increase in ionisation energy occurs between the electron in the outer shell and the next electron in the inner shell. The number of electrons before the large jump gives the group number.

Explain how successive ionisation energies provide evidence for the electron shells in sodium.

There are large differences/increases in ionisation energy between the 1st and 2nd ionisation energies, and between the 9th and 10th ionisation energies. These large increases show the presence of different shells/energy levels.

Explain how successive ionisation energies provide evidence that magnesium is in Group 2.

There is a large jump/increase between the 2nd and 3rd ionisation energies. The 3rd electron is removed from a new shell closer to the nucleus. This shows there are 2 electrons in the outer shell, so magnesium is in Group 2.

Explain how successive ionisation energies can be used to show the sub-shell structure of atoms.

The first ionisation energies of magnesium and aluminium show that the outer electron in magnesium is in the 3s sub-shell and the outer electron in aluminium is in the 3p sub-shell. The 3p sub-shell is higher in energy than the 3s sub-shell, so less energy is needed to remove the 3p electron.

Identify the element from successive ionisation energies.

The largest increase occurs between the 4th and 5th ionisation energies. The 5th electron is in a new shell closer to the nucleus with less shielding. The element has 4 electrons in its outer shell. The element is in Group 4. In Period 3, the element is silicon (Si).

Identify the element from successive ionisation energies (Group 2).

The largest increase occurs between the 2nd and 3rd ionisation energies. The 3rd electron is in a new shell closer to the nucleus with less shielding. The element has 2 electrons in its outer shell. The element is in Group 2.

Define periodicity.

The repeating pattern of trends in physical and chemical properties across different periods.

State the order of elements in the Periodic Table.

Elements are arranged in order of increasing atomic number (number of protons in the nucleus).