Chemistry Summaries

The Periodic Table

Periodicity refers to the periodic trends - the trends of the elements' properties that can be seen on the periodic table.

Atomic Radius

The radius of an atom is the distance from the center of its nucleus to the edge of its valence shell. Since this is incredibly difficult to measure directly, scientists take the distance between the nuclei of two atoms of the same element bonded to each other, and divide that distance by two. For noble gases, they do the same, except noble gas atoms don't bond to each other and are separated by a very small amount of space, so using this method gives an answer that is slightly larger than the real answer.

Since atoms in different periods have a different number of shells $atoms in the 1st period have one shell, atoms in the 2nd period have 2 shells, etc.$, the atomic radius increases as you go down a group in the Table. For example, Lithium $Li$ has a larger atomic radius than Hydrogen $H$, Sodium $Na$ has a larger atomic radius than Lithium, and Potassium $K$ has a larger atomic radius than Sodium.

The more protons there are in the nucleus, the more strongly the atom's nucleus can pull on the electrons, and the closer the electron shells end up being to the nucleus. Therefore, the atomic radius decreases as you go from left to right in a period of the Table.

Combining these two rules, Helium $He$, at the top-right corner of the Periodic Table, has the smallest atomic radius, and Francium $Fr$, at the bottom-left corner of the Periodic Table, has the largest atomic radius.

Ionic Radius

When atoms form ions, they lose or gain valence electrons. But since they don't lose protons, the remaining electrons are still being attracted to the nucleus with the same force, even though that force is acting on fewer or more of them. As a result, the electrons in the valence shell are pulled closer or pushed farther, and the atom's radius changes. This is the atom's ionic radius.

Ionic radii are a bit more difficult than atomic radii. Since elements beyond the 3rd period of the Periodic Table can form more than one stable ion, their ionic radii won't be covered here.

Since atoms in different periods have a different number of shells, the ionic radius increases as you go down a group in the Table.

Metals form cations, which means that they lose their valence shell electrons in ionization. This causes the metal cations to have one less electron shell than their neutral counterparts. However, since the number of protons doesn't change in ionization, elements further along a period $further to the right$ have more protons than the ones before them, even when they're ionized. The more protons there are in the nucleus, the more strongly the atom's nucleus can pull on the electrons, and the closer the electron shells end up being to the nucleus. Therefore, the ionic radius decreases as you go from left to right in a period of the Table, for the metals.

Non-metals, however, form anions, meaning that they gain electrons to fill their valence shell in ionization. Non-metal anions thus have the same number of electron shells as their neutral counterparts. They also have the same number of protons, but since these protons are now pulling on more electrons, they pull less strongly on each of them, and the electron shells end up being farther from the nucleus than in their neutral counterparts. This causes a "split" at the 4th group $the Carbon group$ of the Periodic Table, where elements after it have a larger ionic radius than elements before it due to not losing a valence shell in ionization. The general pattern of the ionic radius decreasing as you go from left to right still holds, there's just a split at the Carbon group of the Periodic Table, where elements after the Carbon group have a larger ionic radius than elements before it due to not losing a valence shell in ionization.

Electrical Conductivity

If an element is a metal, it conducts electricity. Otherwise, it can't conduct electricity.

Melting Points

The melting point of an element is determined by the strength of the intermolecular forces of that substance. It thus follows that metals have higher melting points than non-metals, since for metals, the metallic bonding is an intermolecular force, while for non-metals, the intermolecular forces are much weaker than the intramolecular bonds, and thus weaker than metallic bonding.

However, elements with giant covalent structures $carbon, silicon, etc. Group 4 elements$, the melting points are even higher than those for metals, as the intramolecular bonds which also function as intermolecular bonds are stronger than they are in metals.

Within metals, and elements with giant covalent structures, elements higher up in the group will have higher melting points, and within the same period, elements with more protons in the nucleus will have higher melting points. In non-metals, the opposite is true.

The difference between the groups is that metals will always have higher melting points than non-metals, and elements with giant covalent structures will generally have higher melting points than metals.