S3.1.3 Electron shielding/effective nuclear charge

Electron shielding

When inner shielding electrons reduce the electrostatic force (attraction) between the outer valence electrons and the nucleus.

Shielding electrons

Electrons in energy levels n=1 to n=3

The difference between the amount of charge required to remove valence electrons versus inner electrons is…

that inner electrons require significantly more energy, whereas valence electrons require less energy.

Electron shielding across a period..

remains constant as shielding electrons don’t change

Electron shielding down a group..

increases as shielding electrons increase

Effective nuclear charge

Net positive charge experienced by valence electrons

Atomic radius across a period…

decreases, because:

Nuclear charge increases

Shielding effect constant

Attraction between nucleus and outer electrons increases

Atomic radius down a group…

increases, because:

number of occupied energy levels increases

Which period has a large increase in ionic radius?

Period 3 from silicon ions to phospide, because:

There is an additional occupied energy level

For first 3 ions in period 3…

Ionic radius decreases, because:

Electron configuration stays the same, but atomic number increases.

For ions of elements nitrogen, oxygen, fluorine in period 2…

Ionic radius decreases, because:

Electron configuration stays the same, but atomic number increases.

Positive ions have ______ radii than atoms

Smaller, because:

Positive ions must lose electrons to obtain a full outer shell due to larger charge difference

Thus, there is a stronger electrostatic force of attraction

Negative ions have ______ radii than atoms

Bigger, because:

Negative ions must gain electrons to obtain a full outer shell

Thus, there is a weaker electrostatic force of attraction

First ionization energy

Energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous ions

Ionization energy’s sign is…

positive, as energy is required to overcome the electrostatic force between the valence electrons and the nucleus

Ionization energy across a period..

increases, because:

Nuclear charge increases, atomic radius decreases across period

Thus, greater electrostatic force of attraction

Ionization energy down a group..

decreases, because:

Number of occupied levels increases down group, increased electron shielding

Thus, weaker electrostatic force of attraction

First electron affinity

Energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions

Second electron affinity

Energy released when one mole of electrons is added to one mole of gaseous 1- ions

Electron affinity _________ when atomic radius is greater and electron shielding is greater

decreases

Second electron affinity values are positive when..

the ion is already negative (due to extra repulsion when adding negative electrons)

The difference in the electron affinities of non-metals and metals is..

Non-metals have more exothermic electron affinities (greater positive value)

Electronegativity

Measure of the attraction of an atom for a bonding pair of electrons

Electronegativity across a period..

increases, because of:

increasing nuclear charge, as this increases attraction force.

Electronegativity down a group..

decreases, because of:

increasing atomic radius

Metallic character

How easily an atom can lose electrons (opposite of electronegativity)

Metallic elements’ ionization energies..

are low, and they tend to lose electrons to form positive ions

Non-metallic elements’ ionization energies..

are high, and they tend to gain electrons to form negative ions

Metallic character down a group..

increases, because of

increased atomic radius (due to more energy levels being occupied) resulting in weaker electrostatic force of attraction.

Metallic character across a period..

decreases, because of:

increased nuclear charge, and smaller atomic radius, resulting in stronger electrostatic force of attraction.