Topic 1.4: The periodic table I

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1.4a The order and electronic configuration of elements
Each of the elements in the periodic table is arranged in order of ascending atomic number (proton number). For example we can see that shown in the form
_{Atomic Number}NAME , the first few elements are; ₁H, ₂He, ₃Li, ₄Be, and ₅B etc. 

A set of rules are used to place electrons into atoms.  
1. The lowest shell/quantum shell/energy level is filled first, i.e. 1st before 2nd before 3rd etc.
2. The lowest sub-shell is filled before a higher one, i.e. s before p before d.
3. Orbitals within a sub-shell hold single electrons before being holding pairs.
4. If an orbital has two electrons they have opposite spins.

Their electron configurations of the first five elements are as follows;
H :- 1s¹ He :- 1s² Li :- 1s² 2s¹ Be :- 1s² 2s² B :- 1s² 2s²p¹
These gradually increase up to larger atoms such as Krypton which has an electron configuration as follows. Kr :- 1s² 2s²2p⁶ 3s²3p⁶3d¹⁰ 4s²4p⁶
The electrons in boxes representation is shown below.

Task 1.4bi
half and full d sub-shells are specially stable so Cu :- 1s² 2s²p⁶ 3s²3p⁶3d¹⁰ 4s¹4p

Note 3d¹⁰ 4s¹ not 3d⁹ 4s² .
Task 1.4bii  Draw the electronic configurations as 1s², 2s² etc and the electrons in boxes diagrams for the following atoms: ₂₃V, ₂₄Cr, ₂₅Mn, ₂₆Fe, ₂₇Co.

1.4b Explaining trends in physical properties in period 3 from sodium to argon in terms of structure and bonding

The trend in melting temperatures and in boiling temperatures is for an increase from sodium to silicon followed by a rapid drop for the rest of the period.
For sodium, magnesium and aluminium the melting and boiling temperatures increase as the metallic bonding involved becomes stronger as more electrons are added to the sea of delocalised electrons.  One for sodium, two for magnesium and three for aluminium.
Silicon has extremely high melting and boiling temperatures as this has a giant atomic structure with strong covalent bonds holding each atom in place.  Sulphur and phosphorus and chlorine have simple molecular structures so have low melting points as only Van der Waals forces hold molecules together.  Argon consists of a simple molecular structure with single atoms, each with few electrons, being held by weak Van der Waals forces.

The trend in electrical conductivity across period 3 is for it to increase from sodium to aluminium, to fall greatly for silicon (a semiconductor) then to drop severely so that the remaining elements should be considered insulators.

The metal elements from sodium to aluminium have increasing numbers of electrons available to put into the delocalised sea of electrons.  More electrons are therefore available to carry current and the conductivity increases.  Silicon has its atoms in a giant structure but it is very hard for it to release electrons to carry current.  However this is just possible.  The other non-metals have no charged particles available to carry current and so are insulators.

The trend in first ionisation energies is for a gradual general increase as shown below:

From left to right across the period the nuclear charge is increasing by 1 from each element to the next.  The new electron which is added is put into the same energy level.  The result is that the outer electrons are held more firmly each time the atomic number increases by one.  For aluminium the electron removed is from the 3p orbital which leaves a complete 3s orbital.  This requires less energy than the removal of the outer electron in magnesium which must come from a complete 3s orbital.  For sulphur there is no increase because the electron being removed is from a paired 3p orbital.  The electron removed from phosphorus comes from a half filled 3p sub-shell which is more stable and is energetically slightly more difficult.

Plot and explain the shape for second ionisation energy against atomic number for period 3 using this excel spreadsheet.