Periodic Table Stuff    

Created by: Travis Doyle

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            Periodic Trends

Atomic Radius

    The atomic radius is the distance from the electron from the proton. The size of the electron cloud increases as the principal quantum number increases. you look down the periodic table, the size of atoms in each group is going to increase. When you look across the periodic table, you see that all the atoms in each group have the same principal quantum number. However, for each element, the positive charge on the nucleus increases by one proton. This means that the outer electron cloud is pulled in a little tighter. One periodic property of atoms is that they tend to decrease in size from left to right across a period of the table. 

OVERALL:  Atomic radii increase as you go down and left on the periodic table.

Ionization Energy

     Ionization energy is the energy needed to remove the most loosely held electron from an atom. Ionization energies are periodic. The ionization energy tends to increase as atomic number increases in any horizontal row or period. In any column or group, there is a gradual decrease in ionization energy as the atomic number increases. Metals typically have a low ionization energy. Nonmetals typically have a high ionization energy. 

OVERALL: Ionization energy increases as you go up and right on the periodic table.

Electron Affinity

    Electron affinity is the attraction of an atom for an electron. Metals have low electron affinities while nonmetals have high electron affinities. The general trend as you go down a column is a decreasing tendency to gain electrons. As you go across a row there is also a trend for a greater attraction for electrons.

OVERALL:  EA (electron affinity) increases as you go up and right on the periodic table.



    Electron affinity is the ability of an atom to attract electrons from a bond.  As the charge in the nucleus increases, the atom is able to attract more electrons towards itself (remember that opposites attract).  Since noble gases do not bond with other elements naturally, they are not included as we consider electronegativity.  

OVERALL:  Electronegativity increases as you go up and right on the periodic table.


Electron Configuration

Electron configuration: Is the way in which electrons are distributed among the various orbitals. 

*** Important Rules/Principles/Theories ***

Hund's Rule: The most stable arrangement of electrons is that with the maximum number of unpaired electrons, all with the same spin direction. (One electron goes in a orbital at a time before doubling up)

Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.  (Only 2 electrons can be in an orbital and they must have opposite spins)

Core electrons: Inner electrons of an atom (all electrons except the valence electrons)

Valence electrons: Outermost electrons of an atom (in highest energy level). These determine the chemical properties of an element.

Paramagnetic: An atom has unpaired electrons in its electron configuration. (Look at its orbital diagram)

Diamagnetic: All electrons in an atom are paired. (Look at its orbital diagram)

Aufbau principle:  Electrons fill the lowest energy level subshell first before moving to the next level.



Writing electron configurations: 3 different ways

1.Orbital box notation- this is the way shown in the orbital diagram below

2.Spectroscopic notation (more commonly used, but the "long way")- this is the was it is written under the electron configuration below on the right.

3. Noble gas core notation:  This is the short way.  Instead of writing out all of the electrons in the configuration, you can write out just the ones since the last noble gas.  Find carbon on the periodic table, and then go backwards until you reach a noble gas.  In this case, it is helium.  We can use a shorthand to indicate all of the electrons that are identical to helium's configuration by putting He in square brackets, and substituting it for those electrons.

Noble Gas Notation:  C = [He]2s22p2   

or maybe Te = [Kr]5s24d105p4



Write the electronic configuration for the following atoms or ions.  Use the box notation, spectroscopic notation and noble gas notation.  Are they paramagnetic or diamagnetic?
For the answers, scroll to the very bottom!

1. H
2. He
3. C
4. Ne
5. Cl-


Protons, Electrons, Neutrons in a atom, ion, or isotope

Step 1 - Gather Information

The first thing you will need to do is find some information about your element. Go to the Periodic Table of Elements and click on your element. If it makes things easier, you can select your element from an alphabetical listing.

Use the Table of Elements to find your element's atomic number and atomic weight. The atomic number is the number located in the upper left corner and the atomic weight is the number located on the bottom, as in this example for krypton:

Krypton's data from the Table of Elements

Step 2 - The Number of Protons is...

The atomic number is the number of protons in an atom of an element. In our example, krypton's atomic number is 36. This tells us that an atom of krypton has 36 protons in its nucleus.

What is cool is that is that every atom of krypton contains 36 protons. If an atom doesn't have 36 protons, it can't be an atom of krypton. Adding or removing protons from the nucleus of an atom creates a different element. For example, removing one proton from an atom of krypton creates an atom of bromine.

Step 3 - The Number of Electrons is...

By definition, atoms have no overall electrical charge. That means that there must be a balance between the positively charged protons and the negatively charged electrons. Atoms must have equal numbers of protons and electrons. In our example, an atom of krypton must contain 36 electrons since it contains 36 protons.

Electrons are arranged around atoms in a special way. If you need to know how the electrons are arranged around an atom, take a look at this page.

An atom can gain or lose electrons, becoming what is known as an ion. An ion is nothing more than an electrically charged atom. Adding or removing electrons from an atom does not change which element it is, just its net charge.

For example, removing an electron from an atom of krypton forms a krypton ion, which is usually written as Kr+. The plus sign means that this is a positively charged ion. It is positively charged because a negatively charged electron was removed from the atom. The 35 remaining electrons were outnumbered by the 36 positively charged protons, resulting in a charge of +1.

Step 4 - The Number of Neutrons is...

The atomic weight is basically a measurement of the total number of particles in an atom's nucleus. In reality, it isn't that clean cut. The atomic weight is actually a weighted average of all of the naturally occurring isotopes of an element relative to the mass of carbon-12. Didn't understand that? Doesn't matter. All you really need to find is something called the mass number. Unfortunately, the mass number isn't listed on the Table of Elements. Happily, to find the most common mass number, all you need to do is round the atomic weight to the nearest whole number. In our example, krypton's mass number is 84 since its atomic weight, 83.80, rounds up to 84.

The mass number is a count of the number of particles in an atom's nucleus. Remember that the nucleus is made up of protons and neutrons. So, if we want, we can write:

Mass Number = (Number of Protons) + (Number of Neutrons)

For krypton, this equation becomes:

84 = (Number of Protons) + (Number of Neutrons)

If we only knew how many protons krypton has, we could figure out how many neutrons it has. Wait a minute... We do know how many protons krypton has! We did that back in Step 2! The atomic number (36) is the number of protons in krypton. Putting this into the equation, we get:

84 = 36 + (Number of Neutrons)

What number added to 36 makes 84? Hopefully, you said 48. That is the number of neutrons in an atom of krypton.

The interesting thing here is that adding or removing neutrons from an atom does not create a different element. Rather, it creates a heavier or lighter version of that element. These different versions are called isotopes and most elements are actually a mixture of different isotopes.

If you could grab atoms of krypton and count the number of neutrons each one had, you would find that most would have 48, others would have 47, some would have 50, some others would have 46, a few would have 44 and a very few would have 42. You would count different numbers of neutrons because krypton is a mixture of six isotopes.

In Summary...

For any element:

Number of Protons = Atomic Number

Number of Electrons = Number of Protons = Atomic Number

Number of Neutrons = Mass Number - Atomic Number





1. H = 1 electron (atomic number from the periodic table = 1)

We will start with the lowest energy subshell, the 1s, which has 1 orbital.  The single electron is placed in this orbital.

Box notation:  We can explicitly show each electron in each orbital.  Each orbital is shown as a box.  Each electron is shown as an arrow.  An arrow point up indicates spin up, and an arrow point down indicates spin down.  Each subshell of orbitals is labeled underneath the grouping of orbitals for that subshell.

Hydrogen box notation

Spectroscopic notation:  The orbital is named, and the number of electrons inside the orbital is shown as a superscript.

Hydrogen atoms are paramagnetic with one unpaired electron.

2.  Helium:
The atomic number is 2, so He has 2 electrons to place.

Spectroscopic Notation:  He = 1s2

Box notation:

Helium box notation

Note that helium is diamagnetic since all of the electrons are paired.  Note, also, that the 1s subshell is now filled since it only contains one orbital and can hold at most 2 electrons.

3.  Carbon:
Atomic number = 6, so C has 6 electrons to place.
The 1s subshell is filled with just 2 electrons.  From the Aufbau diagram above, we see that the next subshell to fill is the 2s.  It will also be filled with 2 electrons, so we will have to move to the 2p subshell for the remaining 2 electrons.

spectroscopic notation:  C = 1s22s22p2

Noble gas core notation:  Instead of writing out all of the electrons in the configuration, we can write out just the ones since the last noble gas.  Find carbon on the periodic table, and then go backwards until you reach a noble gas.  In this case, it is helium.  We can use a shorthand to indicate all of the electrons that are identical to helium's configuration by putting He in square brackets, and substituting it for those electrons.

Noble Gas Core Notation:  C = [He]2s22p2

In this case, we did not really save any effort, but for much bigger atoms, the noble gas core notation can be very convenient.

Box notation: Each s subshell has one orbital, but the p subshell has 3 orbitals.  Hund's rule tells us that we have to put the 2p
electrons in separate orbitals since there is room to do so.

Carbon box notation

Thus, carbon is paramagnetic with two unpaired electrons.

4.  Neon:
Atomic number = 10, so there are 10 electrons to place.

Spectroscopic Notation:  Ne = 1s22s22p6

Noble gas core notation: [Ne]

Box notation:

Neon box notation

Since all of neon's electrons are paired, it is diamagnetic.  The six electrons in the p subshell completely fill it.  If you needed to place 11 electrons (for sodium) you would have to go to the next subshell, 3s.

5. Cl-
Atomic number = 17 for Cl;  Add 1 for the anion = 18 electrons in Cl-

Spectroscopic Notation:  1s22s22p63s23p6

Noble gas core notation: [Ar]

Box notation:

Chloride ion box notation


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