What happens to the attraction between the nucleus and the outermost electrons as you move down a group?

The size of an atom is determined by the electrostatic attraction between the positive nucleus and the valence electrons. The force of this attraction is proportional to the size of charge and inversely proportional to the distance between the charges.

As we go down a group the distance between the nucleus and valence electrons increases, while the charges involved stay the same. This obviously has the effect of weakening the force of attraction and the valence electron is held with less force. As a consequence, the valence electrons orbit further away from the nucleus and this gives the atom a greater size.

Since the reducing strength of the atom depends on its ability to give away its valence electron it makes sense that the reducing strength of an atom increases as we move down a group. Since the force of attraction of the electron to the nucleus decreases as we move down so should the ability for the atom to give away its electron increase.

The first ionisation energy and electronegativity also decrease down a group as a result of the weakening force of attraction.

Across a period, the distance between the nucleus and the valence electrons remains constant but the effective core charge increases. As a result the force of attraction between the nucleus and the valence electrons increases across a period. Look at the animation above. Since the force of attraction increases across a period the trends noticed down a group are reversed.

The columns (downwards) of a periodic table are called groups. The rows (across) of a periodic table are called periods.

​Periodic trends can be seen in atoms and ions as you move across a period and down a group. The trends investigated will include Atomic and Ionic radii (distance from the centre of the nucleus to the valence energy level), electronegativity and the 1st ionization energy

Atomic Radii is affected by two main factors : (1) Nuclear charge (number of protons) : The stronger the ‘pull’ the protons have to the electrons with electrostatic attraction, then the smaller the size of the atom radii (2) Number of energy levels: The greater the number of energy levels, the larger the atomic radii. The internal  energy levels “shield” and reduce electrostatic attraction of the valence electrons to the protons.​Addition of another electron does not result in a fractional decrease in the electrostatic attraction to any given electron, but it does increase the electron-electron repulsion, so an overall decrease in Net attractive force. [The most significant factor is nuclear charge – atoms are neutral and as the number of protons increases, the number of electrons increases at the same rate. The increased numbers of electrons in the valence shell, also have an increased number of equivalent protons, and they will be pulled in tighter and therefore show a smaller atomic radii.]

Across the period the Atomic radii decreases​e.g. Li → Ne in period 2 As the nuclear charge increases across the period, so does the electrostatic attraction and so outer electrons are pulled closer to nucleus. The electron repulsion is balanced by the nuclear charge attractions, and as the nuclear charge gets larger, so the electrons get closer together.  The net attractive electrostatic attraction is increased as the nuclear charge increases.

Down the group the Atomic radii increases​e.g. Li → Fr in group one 

Electrons are being added to successive energy levels and both charge on nucleus and electron repulsion increase in step to “cancel each other out”. However, successive energy levels are further from the nucleus therefore, there is a subtle increase in atomic radii and an overall decrease in Net electrostatic attraction, with an increase the electron-electron repulsion.

Cations (metal ions) are smaller in radii than their atomsThe outside energy level of electrons are removed but the nuclear charge (number of protons) remains the same creating smaller radii than the atom. There is an increase in Net electrostatic attraction. Anions (non-metal ions)  are larger in radii than their atoms Extra electrons are added to the outside valance shell that have to be accommodated for and there is no change to the nuclear charge. Electron-electron repulsion spreads the electrons out further creating larger radii than the atom. There is a decrease in Net electrostatic attraction.

Cations i.e. Na to Na+ The inter-electron repulsion experienced by the electron cloud of the cation is less than the neutral atom (less electrons than protons) , and since both species have the same amount of nuclear charge, the net electrostatic attractive force on the electron cloud in the cation is greater than the neutral atom resulting in a smaller cation size.

​Anions i.e. Cl to Cl- The inter-electron repulsion experienced by the electron cloud of the cation is greater than the neutral atom (more electrons than protons), and since both species have the same amount of nuclear charge, the net electrostatic attractive force on the electron cloud in the cation is less than the neutral atom resulting in a larger anion size.

​The first ionisation energy (1st I.E.) is the energy required to remove one mole of electrons from the outside valance shell of 1 mole of atoms in a gaseous state.

An Ionisation equation can be written as:   (M representing an atom)M     →      M+    +    e-            ∆H = + kJ    as this is always an endothermic reactionIf the ionisation energy is high, that means it requires a lot of energy to remove the outermost electron. If the ionisation energy is low, that means it takes only a small amount of energy to remove the outermost electron.Ionisation energy is affected by two factors:

(1) Nuclear charge: As NC increases, there is a stronger pull to the electrons by electrostatic attraction → increased 1st I.E. with increased NC

​In order to remove an electron from an atom you need to overcome the nuclear attraction of its protons holding it around the nucleus. This can be shown by the mud that the car is stuck in. The more mud (nuclear attraction) the more energy to remove the car (electron). However, other electrons in the atom are repelling the electron to be removed – so the more people pushing the car (electron repulsion) the easier it is to extract the car (electron)

Across a period the 1st ionisation energy increases
As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove an electron from the outermost energy level and that means there is a higher ionisation energy. As you go across the Periodic Table, nuclear charge is the most important consideration. Therefore, going across the Periodic Table, there will be a trend of increasing 1st ionisation energy, because of the increasing nuclear charge.

Down a group the 1st ionisation energy decreasesGoing down a group in the Periodic Table, the effect of increased nuclear charge is weighed against the effect of increased electron repulsion, and the number of energy levels becomes the predominant factor. With more energy levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so strongly attracted to the nucleus, and therefore, there is a reduction in net electrostatic attraction. Thus, the ionisation energy of the elements decreases as you go down a group in the Periodic Table, because it is easier to remove the electrons. The more stable elements have higher ionisation energies.

Electronegativity is the tendency of an atom to attract bonding electrons from another atom. Higher electronegativity values mean a higher tendency to attract electrons. Atoms with high E.N. are strong oxidants (gain electrons).Electronegativity is affected by two factors:

1. Nuclear charge: As an atom's nuclear charge increases, there is a stronger pull on electrons of another atom by electrostatic attraction.

​Bond types between atoms can depend on the electronegativity of the atoms. Rather than discrete categories, molecules fall along a continuum

​If there is little difference in electronegativity between two atoms then they tend to form a covalent bond with no polarity difference. A greater electronegativity difference creates a polar bond with uneven “sharing” of valance electrons.
If the electronegativity is even greater then there will be a complete transfer of electron from one atom (Metal) to another atom (non-metal) and ions will form that are held together with an ionic bond.

Across a Period the electronegativity increases
e.g. Li → Ne: The atoms have increased “pulling power” as the nuclear charge is increasing. Electrons are held tighter to the nucleus and there is a greater net electrostatic attraction. This allows another atom to be closer and it has a stronger attraction to electrons from that atom, so electronegativity increases.

Down a group the electronegativity decreases
e.g. Li → Fr   Both nuclear attraction and electron repulsion increase in step, but with an overall decrease in Net electrostatic attraction. However, as successive shells increase atomic radii, then the electrostatic attraction of the nucleus to other atoms’ electrons decreases, so atoms have less electronegativity as you move down a group.