Periodicity Type of Element

Pirate Chemistry 

Periodicity 

The Periodic Table as it exists now gives a wealth of information to Chemists. Knowing how 

to interpret this information can make your life in Chemistry much easier. 

Type of Element 

If we examine the table, we can see that it is broadly divided into three main sections by the 

heavy, black stairstep line. Those elements to the left of the stairstep line are metals. You can 

see by the table that the vast majority of all elements are metals. 

Metals generally: 

Are shiny 

Conduct electricity and heat 

Are malleable (can be hammered into thin sheets or foils) 

Are ductile (can be drawn into wires) 

Those elements to the right of the stairstep line are non-metals. 

Non-metals generally: 

Are dull-looking 

Don’t conduct heat or electricity well 

Are brittle as solids or are gases 

Those elements along the stairstep line are metalloids or semi-metals. They often have characteristics 

of both metals and non-metals. For instance, Silicon is shiny and conducts electricity 

like a metal but is brittle when hit with a hammer. 

Non-metals 

Metals 

Metalloids 

All text copyright Chris Smith 2009. All pictures obtained from internet and are copyright of their owners but assumed 

to be public accessible. If you are the owner of a picture and want it removed, email csmith@d211.org, and 

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Sections of the Table 

You can also see that the table is broken into four major sections: the two tall columns on the 

left; the six tall columns on the right; the middle ten short columns; and the fourteen boxes removed 

from the table and placed at the bottom. Broadly speaking, the tall columns are called 

Main Group elements. They are columns 1, 2, and 13-18 or columns 1A through 8A (the A 

meaning a main group). These elements tend to follow trends and expectations fairly closely. 

The middle ten columns, 3 through 12 or 1B through 8B are called Transition Metals. Their 

chemistry can get a little tricky to predict so at this level we often ignore them. The two long 

rows at the bottom are called the Lanthanide and Actinide series. These elements are the 

strangest of all and we will be doing very little with them. The only reason these elements are 

at the bottom is so that the periodic table fits nicely on a sheet of paper. Examine the table at 

the bottom of the page for what the table really looks like! 

Main Group 

Transition Metals 

This is the extended periodic table. The Lanthanide 

and Actinide series are properly placed but it 

makes the table so wide, that it is often though to 

fit on a page! 



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Groups or Families 

Look at any periodic table and you’ll notice that each vertical column on the table is labeled 

with a number and often a letter. This is to designate a group or family. A group or family on 

the periodic table are elements that are similar to each other. You probably resemble your family 

members in your physical and/or behavioral tendencies. This is not to say you are identical 

to your family, every person is unique, but whether you want to admit it or not, you are probably 

much like your family. The elements on the periodic table are the same way. Any element 

in the same column will most likely share many of the same qualities of the other elements in 

that column. 

Some families are so similar that they are given special names. These are the Noble Gases, the 

Alkali Metals, the Alkaline Earth Metals, and the Halogens. The Noble Gases (group 8A or 

18) are the elements in the very last column of the periodic table. They are called such because 

in ancient times the Noble families did not mix with the commoners. The Noble Gases do not 

react with any other element (unless under very unique circumstances) and thus cannot make 

compounds. The Alkali Metals (group 1A or 1) are the elements in the first column of the periodic 

table (Li, Na, K, Rb, Cs, Fr). They are all soft, shiny metals that react violently with water 

to create a basic or alkali solution. The Alkaline Earth Metals (group 2A or 2) are the second 

column (Be, Mg, Ca, Sr, Ba). They are all metals that react with water to form an alkaline 

(basic) solution. They have a different valence than the alkali metals which sets them apart 

from them but that will be covered later. The Halogens (group 7A or 17) are the second to last 

column (F, Cl, Br, I). Halogen in Greek means salt forming and considering the most famous 

salt is table salt, sodium chloride (NaCl), it is easy to see why the family is named as such. 

Alkali Metals 

Noble Gases 

Alkaline Earth Metals 

Halogens 



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Periods and Trends 

While the vertical columns on the table are called groups or families, the rows on the table are 

called periods. Each period represents a different energy level or shell of the electrons in the 

atom. Across each period from left to right there are a number of different trends of a variety of 

physical and chemical properties of the elements. We will focus only on three. 

Atomic Radius 

Each atom is basically a tiny little sphere. The nucleus is at the center and the electrons orbit 

around the nucleus. Electrons from one atom repel the electrons of another atom (due to their 

same charge) and thus the size of the atom depends on the distance the outermost or valence 

electrons orbit. This distance from the nucleus to the valence electrons is called the atomic radius 

and it is different for each atom. 

Valence 

electrons 

Atomic Radius 

+ 

Nucleus 



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http://www.avon-chemistry.com/p_table_atomic_rad.jpg 


The sizes of the atoms are not random, though. There is an obvious trend if you examine the 

picture below: 



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Let’s explain the trends. As we go down a column the atoms get bigger. There are two reasons 

for this. First, the valence electrons are getting further and further away from the nucleus at 

successively larger energy (n) levels. 

n=1 

n=2 

n=3 

n=4 



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Second, as the valence electrons get further and further away, inner electrons prevent the positive 

nucleus from attracting the outer electrons in a process called shielding. The nucleus cannot 

exert as strong of a force on an outer electron when the inner electrons are getting in the 

way. 

Here the nucleus can efficiently attract the electrons as there is 

nothing between the nucleus and the them 

n=1 

Here the nucleus has three levels electrons in-between itself and 

the outermost (n=4) electrons. Levels 1-3 get in the way or 

“shield” the outermost electrons which reduces the ability of the 

nucleus to attract the electrons which allows them to fly further 

away and thus the size of the atom increases. 

n=4 



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As you go from left to right across a row, however, the atoms get smaller. Why is that? 

First, because they are all in the same row, that means they are all at the same energy level. 

This means that each of these atoms have the same amount of shielding as each other. Second, 

as you go left to right the number of protons in the nucleus increases. This means that the nucleus 

has a greater amount of positive charge to attract the negative electrons. As you go down 

a column, the number of protons also increases but the shielding effect is obviously a greater 

influence which is why the size increases top to bottom. 

Ionic Size 

As atoms become ions their size changes in a predictable way. As metals lose electrons they 

drop to lower or inner energy levels and get a larger proton (positive) to electron (negative) ratio. 

This means that the positive effect of the nucleus is larger and therefore able to pull in the 

outer electrons more, decreasing the size of the ion. The opposite is true for non-metals. They 

gain electrons and thus effectively have fewer protons than electrons and therefore the nucleus 

cannot pull the valence electrons in as strongly and we see that the ions get bigger. 

As the ions get more positive, they get smaller. As the ions get more negative, they get bigger. 



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http://visual.merriam-webster.com/images/science/chemistry/matter/atom.jpg 


Chart of Ionic Sizes 

Electronegativity 

Electronegativity is loosely defined as the relative attraction an atom has for electrons. For reasons 

we will get into later, metals want to give away electrons while non-metals want to gain 

electrons. Thus, metals tend to have low electronegativity values while non-metals tend to 

have high values. This is represented by the following table: 

You can see that the highest value is in the upper right, Fluorine, while the least is the lower 

left, Cesium. Generally speaking, electronegativity increases from left to right and decreases 

from top to bottom. Since Noble gases don’t react with anything there is no electron exchange 

so they have no electronegativities and are left off the chart. 



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Ionization Energy 

Ionization Energy is the amount of energy needed to remove an electron from an atom. Since 

non-metals want electrons, they tend to have high ionization energies. Since metals want to 

give up electrons, their ionization energies tend to be low. It is harder to remove an electron 

from a smaller atom due to the nucleus’ greater attraction for it. Thus, we see that ionization 

energy follows an inverse relationship to size. Larger atoms have valence electrons that are 

more shielded and thus easier to “pick off”. With a few exceptions, the ionization energies increase 

from left to right and decrease from top to bottom. 

It is worth noting, however, that there are exceptions to this rule. We see that the B and O columns 

are less than they should be. Let’s see that in another form: 



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B 

O 

Al 

S 

http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Periodic_Trends 

We can explain this phenomenon easily by remembering Hund’s Rule: the most stable arrangement 

of electrons in an atom is when degenerate orbitals are half-filled spinning the same direction 

or totally filled. Let’s examine the orbital box notation of B and Al: 

B 

Al 

[He] 2s 2 2p 1 

[Ne] 3s 2 3p 1 

s 

p 

Each of these atoms has one lone electron in three 

possible orbitals. This is very unstable and thus B 

and Al will more easily lose this electron which 

results in a lower first-ionization energy. 

s p 

B +1 

Al +1 [Ne] 3s 2 

s 

p 

O [He] 2s 2 2p 4 

S [Ne] 3s 2 3p 4 s p 

O +1 

[He] 2s 2 2p 3 

[Ne] 3s 2 3p 3 

S +1 [He] 2s 2 

Here, the fourth electron is violating Hund’s Rule. 

If the atom gets rid of that electron, then it will 

have 3 identical orbitals all half-filled, spinning in 

the same direction. Thus, the first ionization energy 

is lower than expected. 



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Additionally, we can explain ionization energy data past the first. IE 1 is the energy needed to 

remove the first electron from an atom. IE 2 is the energy needed to remove the second and so 

on. If we examine data for different atoms we see a trend: 

Element IE 1 IE 2 IE 3 IE 4 

Na 496 4562 6912 9543 

Why is there such a big jump after IE 1 ? Examine the spectroscopic notation: 

Na: [Ne] 3s 1 

Then remove the first electron with 496 kJ/mol 

Na +1 : [Ne] 

Now sodium has a Noble gas configuration and is stable. To remove another electron 

would violate this and so the IE 2 is much higher. 

What about Magnesium? 


Mg 738 1451 7733 10,540 

The big jump here is after IE 2 . Why? Again, the spectroscopic notation: 

Mg: [Ne] 3s 2 

Then remove the first electron with 738 kJ/mol 

Mg +1 : [Ne] 3s 1 

Mg +1 does not have a Noble gas configuration yet. Remove electron 2 with 1451 kJ/mol 

(IE 2 is higher due to the electrons now being pulled in closer due to the positive charge 

but there still hasn’t been a huge jump yet). 

Mg +2 : [Ne] 

NOW magnesium has a Noble gas configuration and is stable. To remove the next electron 

would require a lot more energy due to the stable configuration. 

You basically see a big jump in Ionization energy after a Noble gas configuration is reached: 


Na 496 4562 6912 9543 

Mg 738 1451 7733 10,540 

Al 578 1817 2745 11,577 



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Electron Affinity 

Electron Affinity is the amount of energy involved when an electron is added to an atom. It is 

basically the opposite of ionization energy. However, the trend is the same as ionization energy 

because metals typically don’t want electrons and have a low electron affinity and non-metals 

want electrons so have a high affinity for electrons. Examine the chart below: 

This looks just like the ionization energy graph before except the exceptions are at different columns. 



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Be 

Mg 

[He] 2s 2 2p 1 

[Ne] 3s 2 3p 1 

s 

p 

Trying to put an electron into this arrangement 

would be very unstable as it would violate Hund’s 

Rule again 

s p 

Be -1 

Mg -1 [Ne] 3s 2 3p 1 

s 

p 

N [He] 2s 2 2p 3 

P [Ne] 3s 2 3p 3 s p 

N -1 

[He] 2s 2 2p 4 

[Ne] 3s 2 3p 4 

P -1 [He] 2s 2 2p 1 

Here, the orbitals are half-filled spinning the same 

direction as so get extra stability from Hund’s 

Rule. Adding in an electron violates Hund’s rule 

and therefore has a lower-than-expected first Electron 

Affinity. 



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Periodicity Type of Element

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