Trends in the Periodic Table
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:
Be B C
Valence e shielded weak
attraction Valence e not shielded strong
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.
2s1 3s1 4s1
core e 10 core e 18 core
e 86 core e
2s1 e weak attraction
7s1 e very weakly attracted
These shielding effects lead to
some very important Periodic Trends
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.
Trends in Atomic Size
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.
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 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.
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.
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.
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.
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 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.
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
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.
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 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
Ions get bigger as you go down because the ion has added energy levels.
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
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.
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
|Nuclear charge increases
|Shielding effect decreases