Questions regarding ionisation energy
1: Why does Ne have a higher first ionisation energy than Ar ?
The electron removed when Ne is ionised is from the n=2 principle quantum shell, whereas in Ar the electron is removed from the n=3 shell. The n=2 shell is nearer to the nucleus and therefore the electron being removed from Ne experiences a stronger attraction to the nucleus, which requires more energy to overcome. There is also less shielding (only the n=1 shell shields this electron from the nucleus, whereas for Ar the n=2 shell also shields the electron being removed from the nuclear charge). These factors outweight the increasing nuclear charge going from Ne to Ar.
2: Why do first ionisation energies increase from Al to Si to P ?
In all three of these atoms the ionised electron is being removed from the n=3 principle quantum shell, so the shielding from n=2 and n=1 shells is the same in each of these atoms. However, going from Al to Si to P there is one more proton in the nucleus each time so the attraction of the nucleus is stronger, and requires more energy to overcome.
3: Why is the first ionisation energy of Na much less than that of Ne ?
The electron removed on ionization of Na is from the n=3 principle quantum shell, whereas in Ne the electron is removed from the n=2 shell. The n=3 shell is further from the nucleus and therefore the electron being removed from Na experiences a weaker attraction from the protons in the nucleus than that being removed from Ne, which requires less energy to overcome. There are also two filled shells shielding the ionized electron in Na whereas only one filled inner shell shields the ionized electron in Ne. These factors outweight the increasing nuclear charge going from Ne to Na.
Questions regarding atomic radius
4: Explain why the radius of the atoms in Period 3 decreases across the period, even though the numbers of protons and neutrons increase.
Across Period 3 all the atoms have three shells of electrons, and in each case these are shielded by the n=1 and n=2 filled inner shells, so across the period the shielding is the same. The number of protons, and hence the nuclear charge, increases across the period and so the attraction between the nucleus and the outer shell electrons increases, pulling the outer shell electrons closer to the nucleus and decreasing the atomic radius.
Questions regarding melting and boiling points
5: The boiling points of sodium, magnesium and aluminium are 883ºC, 1107ºC, and 2467ºC respectively. Explain this trend.
The melting points of these elements reflect the strength of the metallic bonding, and the energy required to break these bonds. The strength of the metallic bonding depends on the charge on the metal ions, and the quantity of delocalised electrons surrounding them. Sodium forms a 1+ ion and donates one delocalised electron per sodium ion, magnesium forms a 2+ ion and donates two delocalised electrons per magnesium ion, and aluminium forms a 3+ ion and donates three delocalised electrons per aluminium ion. It therefore requires the least energy to boil sodium and the most to boil aluminium.
6: Silicon and phosphorus are both covalently-bonded solid elements at room temperature. The melting point of silicon is 1410ºC while that of phosphorus is 44ºC. Explain this difference.
While they are both covalently bonded, silicon forms a giant covalent lattice structure and phosphorus has a simple molecular structure. To melt silicon requires the strong covalent bonds throughout the lattice to be broken, which requires a lot of energy resulting in a very high melting point. To melt phosphorus only requires the weak London forces between the molecules to be overcome, and this requires much less energy.
7: The boiling points of the Period 2 elements nitrogen and fluorine are -195ºC and -188ºC respectively. The corresponding elements in Period 3 do not show the same increase, phosphorus boils at 281ºC and chlorine at -35ºC. Explain why this non-periodic behaviour occurs.
Nitrogen and fluorine are both simple molecules, N2 and F2, and since fluorine has four more electrons than nitrogen, the London forces in fluorine are stronger requiring more energy to overcome. With phosphorus and chlorine, although a chlorine atom has more electrons than a phosphorus atom, the phosphorus molecule is P4 whereas the chlorine molecule is Cl2 and hence the phosphorus molecule has more electrons and therefore stronger London forces and a higher melting point than chlorine.
Other periodic properties – electrical conductivity
8: Which of the metals potassium or calcium has the greater electrical conductivity. Explain your answer.
Calcium has a higher electrical conductivity than potassium. The electrical conductivity in metals depends on the number of delocalised electrons per metal ion. Calcium has 2 delocalised electrons per calcium ion, whereas potassium has 1 delocalised electron per potassium ion.
9: Explain why the electrical conductivity of aluminium is much higher than than of silicon even though and aluminium atom as 3 outer-shell electrons and a silicon atom has 4 outer-shell electrons.
Aluminium donates its three outer shell electrons into a sea of delocalised electrons as part of its giant metallic structure. These delocalised electrons are free to move and carry electrical charge, making aluminium a good conductor of electricity. The four outer shell electrons in silicon are all shared with other silicon atoms as part of bonding pairs throughout a giant covalent lattice. There are therefore no delocalised electrons in silicon which can act as charge carriers.