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Revision as of 09:44, 7 January 2008

Basics of the Periodic Table

Here is what the most common form of the periodic table looks like:

You don't have to memorize the whole thing. However, you should know that:

• a row of the periodic table is called a period.
  • They correspond to the number of electron 'shells' in an atom of an element
• the columns are called 'groups'.
  • They correspond to the number of electrons in the outer 'shell'.
  • Because they have the same number of valence electrons, elements in the same group have similar physical and chemical properties.
• You should be also familiar with the names of groups 1, 7, and 8, which are shown at the top of the table above.

Trends in the Periodic table

Atomic Radius

Atomic radius is, big surprise, the radius of the atom!

Atomic radius increases:

• As you descend a group, since atoms with more electron 'shells' are bigger.
• As you move left along a period, since atoms with more protons but the same amount of energy shells pull their electrons in closer, making them smaller.

So atomic radius increases towards the lower-left corner of the periodic table:←↓.

Electronegativity

Electronegativity has to do with the strength with which an atom pulls electrons towards it.

Electronegativity increases whenever atomic radius decreases, since the smaller an atom is, the stronger the

So electronegativity increases towards the upper-right corner of the periodic table:↑→.

Ionic Radius

This is a tricky one. Ionic radius generally follows the same trend as atomic radius, but there is an important distinction between cations and anions.

• Cations are positive ions: atoms which have lost one or more electrons.
  • Elements in groups 1-3 tend to form cations. For example, Mg forms Mg2+ ions to gain a full outer 'shell' (8 valence electrons).
  • There are more protons than electrons in a cation, so the ionic radius is smaller than that of the corresponding atomic radius.
• Anions are negative ions: atoms which have gained one or more electrons.
  • Elements in groups 6-8 tend to form cations. For example, S forms S2- ions to expose a full outer 'shell' (8 valence electrons).
  • There are more electrons than protons in an anion, so the ionic radius of an anion is greater than its atomic radius.

Therefore, ionic radius follows the same trend as atomic radius (increasing towards the lower-left corner of the Table), except for the fact that anions (those in groups 6-8, generally) have a larger ionic radius than cations.

• Groups 6-8 have larger ionic radii than groups 1-3.
• Within groups 1-3: ←↓
• Within groups 6-8: ←↓

Melting point

Melting point depends on the type of bonding between atoms. The stronger the bond, the higher the melting point.

Reactivity (within a group)

While elements in a group do share similar chemical properties, reacting similarly with other substances, different elements within a group can have different reactivities. The IB expects you to know the following:

Within Group 1

When an alkali metal reacts, it loses its outer electron, leaving it with a full shell of valence electrons.

Further down the group, this extra electron is at a higher energy level to begin with, so it is easier to remove.

Therefore, Reactivity increases with period in Group 1 ↓

This is why potassium reacts more readily than lithium, for example.

Within Group 7

When a halogen reacts, it gains one electron, leaving it with a full shell of valence electrons.

Further down the group, this extra electron is at a higher energy level, so there is less force pulling it into the atom.

Therefore, Reactivity decreases with period in Group 7 ↑

This is why bromine reacts more readily than fluorine, for example.

This is easy to remember since it's the opposite of what happens in Group 1!

Within Group 8

None of the noble gases tend to react with anything. Hihi!