Module 3: Periodicity

Periodicity: a repeating trend of properties across different periods with increasing atomic number

First Ionisation Energy: the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions

Giant metallic lattice: a three dimensional structure of metal atoms held in place by strong metallic bonds

Metallic bonding: the attraction between positive metal ions and delocalised electrons

Delocalised electrons: outer shell electrons that are shared between more than two atoms and are free to move

Giant covalent lattice: a three dimensional network of atoms held in place by strong covalent bonds

Elements are arranged:

• by order of increasing atomic number, i.e. increasing number of protons 
• in horizontal rows called periods, each period repeats trends in physical and chemical properties of elements
• in vertical columns called groups, element in the same group have similar chemical properties, i.e. react in similar ways as their atoms have the same number of outer shell electrons

Some Patterns:

• Across a period elements change from metals to non-metals: horizontal trend
• In group 4, carbon at the head of the group is a non-metal but lead at the bottom of the group is a metal: vertical trend

Patterns in electron configuration:

As you move across a period, atomic number increases by one. The number of electrons therefore increases by one, which is reflected in the electron configuration. From periods 2 and 3:

                ⁴Be 1s²2s^{2                                     12}Mg 1s²2s²2p⁶3s²

                 ⁵B 1s²2s²2p^{1                                   13}Al  1s²2s²2p⁶3s²3p¹                              

                 ⁶C 1s²2s²2p^{2                                   14}Si  1s²2s²2p⁶3s²3p²   

Patterns in atomic radius:

• atomic radius decreases acrosses a period because as nuclear charge increases there is a greater attraction between electrons in the same outer shell and the nucleus
• atomic radius increases down a group because outer shells are further from the nucleus so are increasingly shielded as more inner shells reduce the attraction between the electrons and the nucleus                          

Factors affecting ionisation energy:

• atomic radius (distance from the nucleus)
• nuclear charge (number of protons in the nucleus)
• electron shielding increasing with the number of inner shells

• General increase across periods 2 and 3 with increasing atomic number is due to increasing strength of attraction between an increasing number of protons in the nucleus when electrons are being added to the same outer shell and the atomic radius is decreasing; increasing nuclear attraction
• Decrease down a group is due to increasing atomic radius so a weaker attraction between the outer electrons and the nucleus as outer electrons are further from the nucleus. Greater shielding of outer electrons from more inner shells also leads to less attraction between them and the nucleus
• Small decrease in 1st I.E. between Be and B, and between Mg and Al, due to a difference in s and p subshell energies
• Small decrease between N and O, and between P and S, due to p orbital repulsion
• Successive ionisation energies can be used to predict the number of electrons in the electron shells of atoms of elements, and so predict the group of an element

• Most metals have high melting and boiling points; metallic bonds are strong and extend throughout the structure so a lot of energy is needed to break the bonds; the strength of metallic bonding increases across a period with increasing nuclear charge

• Metals conduct electricity in the solid and liquid states; delocalised electrons are free to move through the giant metallic lattice when a voltage is applied

• Metals are insoluble in solvents; metal atoms are unable to form new attractions with solvent molecules

• All giant covalent structures are solid at rtp with very high melting and boiling points; covalent bonds are strong so a lot of energy is needed to break them

• Giant covalent structures do not conduct electricity in any state; atoms are neutral so cannot conduct electricity (graphite is an exception)

• Giant covalent structures do not dissolve in any solvent; there are no new attractions possible between the atoms and solvent molecules

• Across groups 1 to 4 there is a general increase in melting point
• There is a sharp decrease in melting point from group 4 to 5
• Groups 5 to 8 have relatively low melting points, bit with an erratic pattern for period 3 in particular

• If strong bonds must be broken there will be high melting points as lots of energy is required to separate particles from each other and change the state from solid to liquid

• If only weak forces need to break there will be low melting points and only a little energy is required to separate particles from each other and change the state

Melting points increase due to increasing strength of the metallic bond in giant metallic structures.

weakest bond: Li < Be: strongest bond

weakest bond: Na < Mg < Al: strongest bond

The metallic bond strength increases with an increasing nuclear charge attracting delocalised electrons in the same outer shell:

₁₁Na⁺  has 11 protons    ₁₂Mg²⁺ has 12 protons           ₁₃Al³⁺  has 13 protons

The giant covalent lattices of group 4 have extremely high melting points as strong covalent bonds extend through the structure and must be broken. Boron (group 3) has a giant covalent structure similar to carbon and so has a very high melting point.