Periodicity

Elements are arranged in increasing atomic number in the periodic table Elements in Groups have similar physical and chemical properties The atoms of elements in a group have similar outer shell electron configurations, resulting in similar chemical properties; Elements in periods showing repeating trends in physical and chemical properties Period 2 = Li, Be, B, C, N, O, F, Ne Period 3 = Na, Mg, Al, Si, S, Cl, Ar. Classification of elements in s, p, d blocks Elements are classified as s, p or d block, according to which orbitals the highest energy electrons are in. 0 0.02 0.04 0.06 0.08 0.1 0.12 0.14 0.16 0.18 Na Mg Al Si P S Cl Ar atomic radius (nm) Atomic radius Atomic radii decrease as you move from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell similar shieding. Exactly the same trend in period 2 N Goalby chemrevise.org Periodicity is a repeating pattern across different periods Various properties such as atomic radius, melting points, boiling points and ionisation energy display periodicity

Ionisation Energies Definition :First ionisation energy The first ionisation energy is Energy needed to remove an electron from each atom in one mole of gaseous atoms This is represented by the equation: H(g)
H+ (g) + eAlways gaseous Remember these definitions very carefully The equation for 1st ionisation energy always follows the same pattern. It does not matter if the atom does not normally form a +1 ion or is not gaseous Factors that affect Ionisation energy There are three main factors 1.The attraction of the nucleus (The more protons in the nucleus the greater the attraction) 2. The distance of the electrons from the nucleus (The bigger the atom the further the outer electrons are from the nucleus and the weaker the attraction to the nucleus) 3. Shielding of the attraction of the nucleus (An electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus) Many questions can be answered by application of these factors The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element. Successive ionisation energies 1 2 3 4 5 6 Ionisation energy No of electrons removed Notice the big jump between 4 and 5. Explanation The fifth electron is in a inner shell closer to the nucleus and therefore attracted much more strongly by the nucleus than the fourth electron. It also does not have any shielding by inner complete shells of electron Why are successive ionisation energies always larger? The second ionisation energy of an element is always bigger than the first ionisation energy. This is because the ion formed, is smaller than the atom and the proton to electron ratio in the 2+ ion is greater than in the 1+ ion. The attraction between nucleus and electron is therefore stronger How are ionisation energies linked to electronic structure? 1 2 3 4 5 Ionisation energy kJ mol-1 590 1150 4940 6480 8120 Here there is a big jump between the 2nd and 3rd ionisations energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy Example: What group must this element be in? The first Ionisation energy of the elements The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity. The pattern in the first ionisation energy gives us useful information about electronic structure You need to carefully learn the patterns Many questions can be answered by application of the 3 factors that control ionisation energy Q. Why has Helium the largest first ionisation energy? A. Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton Q. Why do first ionisation energies decrease down a group? A. As one goes down a group, the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller Q. Why is there a general increase in first ionisation energy across a period? A. As one goes across a period the electrons are being added to the same shell which has the same distance from the nucleus and same shielding effect. The number of protons increases, however, making the effective attraction of the nucleus greater. Q. Why has Na a much lower first ionisation energy than Neon? This is because Na will have its outer electron in a 3s shell further from the nucleus and is more shielded. So Na’s outer electron is easier to remove and has a lower ionisation energy. Q. Why is there a small drop from Mg to Al? Al is starting to fill a 3p sub shell, whereas Mg has its outer electrons in the 3s sub shell. The electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons Learn carefully the explanations for these two small drops as they are different to the usual factors Q. Why is there a small drop from P to S? 3s 3p With sulphur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill the first 3p orbital. When the second electron is added to a 3p orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove. 3s 3p Two electrons of opposite spin in the same orbital

Definition: A metallic bond is the electrostatic force of attraction between the positive metal ions and the delocalised electrons The three main factors that affect the strength of a metallic bond are: 1. Number of protons/ Strength of nuclear attraction. The more protons the stronger the bond 2. Number of delocalised electrons per atom (the outer shell electrons are delocalised) The more delocalised electrons the stronger the bond 3. Size of ion. The smaller the ion, the stronger the bond. Metallic bonding Example Mg has stronger metallic bonding than Na and hence a higher melting point. The Metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds.

Macromolecular: diamond Tetrahedral arrangement of carbon atoms. 4 covalent bonds per atom Macromolecular: Graphite Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th outer electron per atom is delocalised. Delocalised electrons between layers. Both these macromolecular structures have very high melting points because of strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds. Bonding Structure Examples Covalent : shared pair of electrons Macromolecular: giant molecular structures. Diamond Graphite Silicon dioxide Silicon Metallic: electrostatic force of attraction between the metal positive ions and the delocalised electrons Magnesium, Sodium Giant metallic (all metals) lattice. Only use the words molecules and intermolecular forces when talking about simple molecular substances Property Macromolecular Giant Metallic boiling and melting points high- because of many strong covalent bonds in macromolecular structure. Take a lot of energy to break the many strong bonds high- strong electrostatic forces between positive ions and sea of delocalised electrons Solubility in water insoluble insoluble conductivity when solid diamond and sand: poor, because electrons can’t move (localised) graphite: good as free delocalised electrons between layers good: delocalised electrons can move through structure conductivity when molten poor (good) general description solids shiny metal Malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another -attractive forces in the lattice are the same whichever ions are adjacent. Melting and boiling points For Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller positive centre also makes the bonding stronger. High energy is needed to break bonds. Si is Macromolecular: many strong covalent bonds between atoms high energy needed to break covalent bonds– very high mp +bp Cl2 (g), S8 (s), P4 (S)- simple Molecular : weak London forces between molecules, so little energy is needed to break them – low mp+ bp S8 has a higher mp than P4 because it has more electrons (S8 =128)(P4=60) so has stronger London forces Ar is monoatomic weak London forces between atoms Similar trend in period 2 Li,Be metallic bonding (high mp) B,C macromolecular (very high mp) N2 ,O2 molecular (gases! Low mp as small London forces) Ne monoatomic gas (very low mp)