Trends in the Periodic Table

Shielding Atomic Size
Ionization Energy Ionic Size
Electron Affinity Electronegativity



Electrons on the outside energy level (valence electrons) have to look through all the inner energy levels to see the nucleus. These inside energy level electrons shield (block) the nuclear force field from the valence electrons. As you go across a row (period) the number of inside electrons is constant. The number of inside energy levels stays the same and these core levels are full. As you go across the row the nuclear charge gets larger (the force field gets bigger) because a proton is added to the nucleus. As you go across the row, a valence electron is added to the valence shell but the valence electrons have the same shielding.

Energy level 2 (Row 2) valence configuration:

                 Li                             Be                          B                     C                                       Ne

                 1e                              2e                           3e                     4e                                     8e  

                 +3                             +4                           +5                     +6                                     +10   


 Valence e shielded weak attraction                               Valence e not shielded strong attraction

Notice the size of the nuclear charge (positive force field) increase from +3 to +10 across the row. The number of inner (core) electrons remains the same (2e) across the row. The inner 2e are between the nucleus and the valence electrons. The blocking strength (shielding effect) of these inner electrons is the same across the row. Across the row the valence electrons will have a greater attraction to the nucleus because of the greater nuclear charge. The shielding becomes less effective across the row 2e can shield 3+ better than 2e can shield +10. The valence electrons in Ne feel a greater positive force field and are more strongly attracted to the nucleus. The valence electron in Li is effectively shielded from the positive force field of the nucleus and is very weakly attracted to the nucleus (metals are born losers – lose valence electrons easily).

What happens to the attraction to valence electrons as you go down a column (family or group)? As you move from one period to the next the valence electrons are being added to a new valence shell (energy level) that is further from the nucleus. As valence shells are filled additional valence electrons must be placed in larger energy levels that are further from the nucleus. These new valence electrons have additional levels of inner shielding electrons (are more effectively shielded from the nuclear charge). As you go down the group the attraction of the valence electrons to the nucleus decreases because these electrons are further from the nucleus and are better shielded from the nuclear force field.

                Li                                Na                             K                                                  Fr

                2s1                               3s1                            4s1                                              7s1

   2 core e                           10 core e                        18 core e                   86 core e

               +3                                +11                                       +19                                                                +87


 2s1 e weak attraction                                                                                7s1 e very weakly attracted

These shielding effects lead to some very important Periodic Trends



Atomic size

Atomic size is measured by using the Atomic Radius = half the distance between two nuclei of a diatomic molecule because the electron cloud doesn’t have a definite edge.


           H               H  

            +                +

             |        |        |         

     Atomic radius

Trends in Atomic Size

Period trends

Within a Period the charge on nucleus increases – the larger charge pulls valence electrons in closer. As you go across a period the radius gets smaller. The smallest atoms in each period (row) will be to the right; the size decreases as you move right along the row.


Group trends

As we go down a group each atom has another energy level. Higher energy level is further away from the nucleus so the atoms get bigger. The largest atoms are found at the bottom of each group. The size of the atom increases as you move down the family.


Ionization Energy

Ionization Energy is the amount of energy required to completely remove an electron from a gaseous atom. The energy required to remove one electron is called the first ionization energy. Removing one electron makes a +1 ion. The second ionization energy is the energy required to remove the second electron. The second ionization energy is always greater than first ionization energy. The third ionization energy is the energy required to remove a third electron and is greater than 1st or 2nd ionization energy.

 What determines Ionization Energy?

1. As the distance from the nucleus increases the attraction to the nucleus decreases. As the attraction to the nucleus decreases the ionization energy decreases. It is easier to remove an electron that is farther from the nucleus.

2. The greater the effective nuclear charge (smaller shielding of valence electrons) the greater the force of attraction to the valence electrons. The stronger the attraction for valence electrons the more energy is needed to remove a valence electron.

3. Filled and half filled orbitals have lower energy, so these configurations would have higher ionization energies, harder to remove an electron from lower energy.


Trends in Ionization Energy

Period trends

All the atoms in the same period have the same valence energy level and the same shielding. The increasing nuclear charge increases the attraction of valence electrons so ionization energy generally increases from left to right. The exceptions at full and ˝ filled orbitals break the pattern because removing an electron from a stable (low energy) configuration will require more energy. Li has a low 1st ionization because 2 core electrons block the +3 nuclear charge more than they could block the +10 nuclear charge in the Ne atom that would have large ionization energy.


Group trends

As you go down a group, first ionization energy decreases because the valence electron is further away and there are more inner electrons causing more shielding. Ne has a lower ionization energy than He even though both are full energy levels because Ne has more shielding and valence electrons are a greater distance from the nucleus.


Text Box: Full sublevel


Text Box: Half filled sublevel



Driving Force – why atoms react

Atoms with a full outer energy level (valence shell) are very low in energy. Noble gases have full valence shells and are therefore very stable elements (not reactive). The Noble gases are like trendsetters while the other elements are wannabes. Atoms behave in ways to achieve the stable noble gas configuration.

2nd Ionization energy - For elements to reach a filled level by removing 2 electrons, the 2nd ionization energy is lower than expected. For example for elements with a ns2 configuration (Alkali earth metals) removing the 2 valence electrons leaves a full energy level and forms a stable +2 ion.

3rd ionization energy using the same logic s2p1 atoms have a low 3rd ionization energy. Atoms in the aluminum family form stable +3 ions with all occupied energy levels full.


Electron Affinity

Electron Affinity is the energy change associated with adding an electron to a gaseous atom. If the atom becomes more stable (electron configuration like the noble gases) there is a loss of energy and the energy change is shown a negative value. If the atom becomes less stable the energy level goes up and the energy change is shown as a positive value. The ability to attract an extra electron depends on the strength of attraction of valence electrons. Strong attraction of valence electrons means a greater the ability to hold an extra electron.

Trends in Electron Affinity

Period trends

Electron affinity is low for metals and high for nonmetals. It is easiest for group 7Ato get and hold an extra electron because it gets them to full energy level. Electron affinity increases from left to right as atoms become smaller, with greater nuclear charge and get closer to a full energy level.

Group trends

Electron affinity decreases as we go down a group because the atoms are getting bigger and the valence electrons are not attracted as strongly to the nucleus.


Ionic Size

Cations form atoms by losing electrons. Cations are smaller that the atom they come from. There are fewer electrons with the same nuclear charge, leading to a greater attraction of the remaining electrons – smaller particle. Metals form cations. Cations of representative elements have noble gas configuration.

Anions form by atoms gaining electrons. Anions are bigger that the atom they come from because there are more electrons being held by the same nuclear charge. Nonmetals form anions. Anions of representative elements have noble gas configuration.

Configuration of Ions

Representative ions always have noble gas configuration. Na is 1s12s22p63s1 . Sodium forms a +1 ion - 1s12s22p6 that is the same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons.


Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.


Trends in Ionic Size

Group trends

Ions get bigger as you go down because the ion has added energy levels.


Period Trends

Across the period nuclear charge increases and energy level changes when the positive ions form. Cations revert back to the previous energy level by losing electrons to attain the previous noble gas configuration., The cations get smaller as you move right because the cations lose one more electron as you move right. Sodium loses the s1 electron, Mg loses the two s2 electrons and the Al loses the three s2p1 electrons. Each cation is isoelectronic (ions have the same number of electrons) with Ne (same electronic configuration). Na1+ has +11 attracting 10e; Mg2+ has +12 attracing 10e and Al3+ has +13 attracting 10e. The ioic size of cations decreases as you move right

Anions gain electrons to form the stable electron configuration. Nitrogen gains 3e(N-3), Oxygen gains 2e (O-2) and Fluorine gains 1e(F-1). Each of the anions will look like Ne(all have 10 electrons). N has +7 attracting 10e, O has +8 attracting 10e and F has +9 attracting 10e. N attraction will be weakest and F attraction will be strongest. The size of the anion decreases as you move right. The anions are bigger than the cations because of the extra energy level in the anion.



Electronegativity is the tendency for an atom to attract electrons to itself while it is chemically combined with another element. It indicates how fair the atom shares its valence electrons. Big electronegativity means the atom pulls valence electrons toward itself. Atoms that have strong attraction to their own valence electrons will be able to attract the valence electrons from other atoms. (Atoms with large negative electron affinity have larger electronegativity.)


Trends in Electronegativity

Period Trend

Metals are at the left have weak attraction for valence electron due to the shielding effect. Metals let their electrons go easily (born losers). Metals have low electronegativity. At the right end are the nonmetals that have strong attraction to valence electrons due to poor shielding. Nonmetals have valence shells that are almost full so they want more electrons and are able to hold them. Nonmetals have high electronegativity. Electronegativity increases as you move across the period.

Group Trend

The further down a group the farther the valence electrons are away from the nucleus and the more inner electrons an atom has to shield the valence electrons. The valence electrons are not held as tightly. As you move down the group the attraction to the valence electrons decreases so the electronegativity will decrease as you move down the group.


Atomic size decreases
Ionization energy increases
Electron affinity increases
Electronegativity increasers
Nuclear charge increases
Shielding effect decreases











































































      Text Box: Atomic radius (size) increases
Ionization energy decreases
Electron affinity decreases
Electronegativity decreases