2.1. Periodicity

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part of a covalent bond in theory, the sum of two covalent radii should equal the covalent bond length between 2 atoms, but in practice the length of the bond depends on the compound the bond is in. Van-der-Waals-Radius O₂ BB Radius van der Waals-Radius 0,062 nm 0,14 nm CI Na Cr Ionic radius Explain the difference between covalent radius and van der Waus radius. → covalent measures the distance from the centre of the covalent bond to the centre of the nucleus, (opposite the bond) vdw measures from the nucleus to the outside of the of the atom where it is attracted to another molecule. Atomic Radius across period 3 0000000 Na Mg Al Si P Size of atoms decreases Atomic radii (pm): 186 160 143 118 110 104 99 Elements: Explain why the atomic radius decreases across period 3. → more protons are added across the period → the nucleur charge increases → draws the electrons closer Summary Questions: 1. Explain why a sulphur atom has a smaller atomic radius than phosphorus (3) → the greater nuclear charge of sulphur draws the electrons closer 2. Explain why the first ionisation energy increases across the third period (3). ->>> electron is removed from the same energy level → nuclear charge increases across the period, making it increasingly difficult to remove an electron 3. Explain why electronegativity increases across the period (3) across the period atomic radius decreases → nuclear charge increases → resulting in a greater attraction for the electrons in a covalent bond. 4. Describe and explain the trend in electronegativity values down group 2. (3) → electronegativity decreases → down the group, since more energy levels are occupied / greater shielding. → nuclear charge increases. Trends in the first ionisation energy. First ionisation energies of Period 3 elements - Explain why the lonisation energy increases across period 3. → electron is removed from the same energy level → nuclear charge increases across the period making it increasingly difficult to remove an electron Questions 1. Mg Mg Al : . : a) Show the electron arrangements (box and arrow diagrams) for Mg and AL. 2 p Element D. D. IS 25 CI Ar 20 35 3p 16 16 16 ✓ b) use these to explain the drop in I.E. the drop in the value for aluminium is because the extra electron has gone into one of the 3p orbitals • this is a higher energy electron than the s electron so takes less energy. Summary questions: first ionisation energy 1. Describe and explain the general trend in first ionisation energies in period 3. -> general increase → due to increasing nuclear charge 2. Explain why the first ionisation energy of aluminium is lower than that of magnesium. AI: 15³ 25² 2p 35² 3p → 1st electron in new 3p orbital so easier to remove than paired e in 35 (Mg) 3. Explain why the first ionisation energy of sulphur is lower than that of phosphorus. → P: 15² 25² 2p 3s² 3p 6 2 → S: 15²³ 25² 2p 35² 3pª Ionisation Energy 0 3rd (n=3) outer shell 4. Preduct the first ionisation energy of the next element potassium. Explain your reasoning. → lower than Na due to increased shielding, therefore easier to remove e- → also low due to it being 1st e² in 45 orbital. 1 2 5. Describe and explain the trend in successive ionisation energies for sodium. 3 4 2nd (n=2) shell electrons 16 6 7 Ionisation Number I 8 3P 30 • first paired e in 3p Spin-pair repulsion = easier to remove 9 10 11 → if you remove one election from an atom at a time, each one requires more energy to remove than the one before 2 electrons in (n-1) shell dosest to nucleus → there is a sharp increase in lonisation energy from the first and secound electrons → this is followed by a gradual increase over the next & electrons 12 and then another jump before the final two elections. Tuesday 11th October 2022 Periodicity Periodicity and properties of group of elements - fill the gaps. • The elements in group 1, 2 and 3 are metals. They have giant Structures. They lose their outer electrons to form ionic (positive ions) compounds. • Silicon in group 4 has 4 electrons in its outer shell, with which it forms 4 covalent bonds. Silicon is classed as a semi-metal. • The elements in group 5, 6 and 7 are non-metals. They either gain electrons to form covalent compounds or Share their outer electrons to form ionic compounds. Argon is in group 0/8 and is a noble gas it has a full outer shell and is unreactive. Trends in melting point Melting and boiling points increase because of the strength of the metalic bonds increases. Na This is because the charge of the ion increates so more electrons join the delocalifed electron sea that holds the metallic Lattice together. ть METALS 371 K Mg Al Si 1156 K 922 K 1380 K 933 K 2740 K SEMI-METAL 1683 K 2628 K P 317 K 553 K SC S Silicon is a giant macromolecular (covalent) structure. very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs. NON-METALS 392K 718 K 172 K Ar 84K 238 K 87K Argon exists as a single atom. There- fore, it's melting point and boiling point are lower than Chlorine. This means VDW forces are weak. The order of melting points of the non-metal molecular Structures depends on the size of the Vow forces. This depends on the number of electrons in the molecule. The more electrons in the molecule, the stronger the VDW forces. Summary Questions: melting points 1. Describe and explain the trend in melting point of the metals sodium and aluminium increase from sodium to aluminium due to larger charge on ions and more delocali red electrons therefore stronger electrostatic forces + metallic bonding. 2. Explain why silicon has such a high melting point →>> forms giant covalent structures. Takes a lot of energy to break covalent bond. 3. What type of structure and bonding are found in elements 15 to 18. → simple covalent → intermollecular forces →vdW. 4. Explain why sulphur Se has a higher melting point than phosphorus P4. →larger molecule therefore larger e cloud therefore stronger vdw therefore higher mp.