Trends on the Periodic Table
Atomic Radius: The size of the atom is determined by a balance between electrostatic
attractions and repulsions. The positive nucleus attracts the negative electrons pulling
them in toward the nucleus. The electrons repel each other. The electron repulsion
however is reduced by the attraction created by the electron spins.
Affect by period
(Slides 1 & 2)
Each time the atomic number increases by on across a period there is an additional
electron and proton added to the atom. This increases both the attraction and
repulsion, but since the repulsion is reduced by spin there is a net increase in attraction
pulling the electrons in closer to the nucleus. Therefore, Atomic Radius decreases going
across a period due to increased attraction between the nucleus and the electrons.
The flatten, u-shaped areas on the graph are due to the filling of the d-orbital which will
push the s-orbital of the next energy level away increasing the size slightly. The
occasional jump in size is created by the start of a new orbital which is slightly further
away from the nucleus.
Affect by family
Each time you go down a row on the periodic table there is an additional energy level
added to the atom. Therefore, Atomic Radius increases going down a family.
(Slide 3)
Bohr’s limits state that there is a maximum of 8 electrons in the outer energy level of an
atom. This is not just a maximum, but a preferred condition. There atomis more stable
when the outer energy level has 8 electron. According to the Octet Rule an atom will
gain, lose or share electrons to obtain a filled outer octet (8 outer electrons). When an
atom gains, or loses electrons the atom becomes charged at which point we refer to it as
an ion.
Ion: charged atom or group of atoms
Ionization: process by which an atom becomes charged.
When an atom becomes an ion it is attempting have an outer most energy level that
contains 8 electrons. This is referred to as a Stable Electron Arrangement. (The first
energy level is stable with 2 electrons)
The normal, uncombined atom is neutral due to equal numbers of p+ and e# of p+ = # of eIn order for atoms to join up to form compounds they need to become charged by either
losing or gaining electrons, filling their outer most energy level.
Valence (Oxidation State): the type of ion formed by an atom. The valence for an
element is given on the periodic table.
Valence electrons: electrons in the outermost orbit of an atom
Positive Ion: atom that has lost electrons.
Example 1: Sodium 11p+
1s2 2s2 2p6 3s1 has 1 outermost electron
To have an outer energy level of 8 electrons it can either gain 7 electrons or lose 1
electron. Since losing 1 electron is a smaller change it will lose one electron to become
1s2 2s2 2p6 3s0
Since the number of protons and electrons is no longer the same this atom now has a
charge.
To find the charge use the formula:
# of p+ - # of e- = Charge
In this example the charge = 11 – 10 = +1. This atom, therefore, has become a +1 ion.
The charge is given as a superscript following the chemical symbol, Na1+
Example 2: Calcium with 20 p+
1s2 2s2 2p6 3s2 3p6 4s2 has 2 outer electrons.
This atom can either gain 6 electrons or lose 2 electrons, since it is easier to lose 2
electrons it will become:
1s2 2s2 2p6 3s2 3p6 4s0
Since 20 – 18 = +2, this atom becomes a +2 ion. Ca2+
Example 3: Scandium with 21p+ 1s2 2s2 2p6 3s2 3p6 4s2 3d1
This atom can either gain 6 electrons or lose 3 electrons (if it loses only 2 electrons the
outer energy level will be left with 9 electrons)
Since losing 3 electrons is a smaller change it will lose 3 electrons to become a +3 ion. (21
– 18 = +3) Sc3+
1s2 2s2 2p6 3s2 3p6 4s0 3d0
Removing electrons requires energy since the electron is being moved away from the
nucleus and the attraction between it and the nucleus is being broken.
Ionization energy: energy required to remove electrons from an atom.
Ionization Energy:
The energy required to remove an electron from the atom.
I.E. increases going across a period due to: (Slide 4)
1) increased attraction between the nucleus and electrons (there are more
electrons and protons attracting each other)
2) decrease in atomic radius (there are more particles attracting and they are
closer together)
The occasional dips on the graph are due to: (Slide 5)
1) starting a new orbital
2) starting to pair electrons in orbitals
I.E. decreases going down a family due the increased distance between the nucleus and
the electrons.
(Slide 6)
Metals therefore have the low Ionization Energies and will tend to lose electrons to
form positive ions. The alkali metals and alkaline earth metals have the lowest
ionization energies making them extremely reactive. Alkali metals and alkaline
earth metals (columns 1 and 2) are so reactive that they are never found uncombined
in nature.
Nobel gases have extremely high Ionization Energies making it very difficult to
make them lose electrons.
Negative Ion: atom that has gained electrons.
In order to gain electrons the atom needs to be able to attract the electrons and hold
onto them. This ability is called electronegativity.
Electronegativity: ability to attract and hold electrons
Example 1: Fluorine 9p+
1s2 2s2 2p5
has 7 outer electrons.
It can either gain 1 electron or lose 7 electrons. (The first energy level is filled with only
2 electrons.) Since it is easier to gain one electron the fluorine atom will become:
1s2 2s2 2p6
Since 9 – 10 = -1 the fluorine atom becomes a -1 ion. F1- Negative ions are named by
changing the ending of the element’s name to –ide. So fluorine becomes fluoride.
To determine the number of valence electrons for a nonmetal subtract the absolute
value of the valence number from 8 (8 - |valence| = valence electrons). I.e. Fluorine
has a valence of -1. The number of valence electrons is 8 – 1 = 7.
Example 2:
Phosphorous has a valence of -3. The number of valence electrons is 8 – 3 = 5.
1s2 2s2 2p6 3s2 3p3
It can gain 3 electrons or lose 5 electrons. Since it is easier for it to gain 3 it will become:
1s2 2s2 2p6 3s2 3p3
This makes it (15 – 18 = -3) a -3 ion. P3- named phosphide.
Electronegativity:
Electronegativity - The ability of an atom to attract and hold onto an electron.
Electronegativity increases going across a period and decreases going down a family.
(Slide 7)
(Slide 8)
The electric field created by the nucleus extends out from
the nucleus beyond the outermost energy level. As the
number of protons in the nucleus increases this field gets
bigger, but the atomic radius decreases leaving a large
field about the atom that can attract passing electrons.
(Slide 9)
Nonmetals have high Electronegativities and Ionization Energies making it easier for
them to gain electrons to form negative ions. The halogens (column 17) are so strong
that they will always be found as a molecule containing 2 atoms when pure. (F2)
Noble gases have complete outer energy levels. They do not need to pick up any
electrons. Helium, Neon and Argon do not have a measurable Electronegativity. This
lack of ability to gain electrons combined with their high Ionization energies make
Noble gases high unreactive (Inert).
Inert – tend to not take part in chemical changes
No compounds have ever been formed with Helium, Neon or Argon. Krypton, Xenon
and Radon will react with Fluorine, Chlorine and Oxygen under extreme conditions of
high temperature and pressure.
Ionic Radius:
(Slide 10)
For a positive ion the ionic radius is smaller than the radius of the neutral atom.
When an atom loses electrons it reduces the repulsions between the electrons orbiting
the atom and reduces the number of occupied energy levels about the nucleus. This
results in a reduction in the radius of the atom.
For a negative ion the ionic radius is larger than the radius of the neutral atom.
When an atom gains electrons, the repulsions, between the electrons orbiting the atom,
increase causing the electrons to move away from each other, resulting in an increase in
the radius of the atom.
The overall pattern is:
(Slide 11)