2.1 Periodicity. Dobereiner Law of Triads:- If you look at the properties and relative atomic masses of 3 elements in group 1:-

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1 2.1 Periodicity The development of the Periodic table: The Periodic Table brings order and a systematic way of looking at the elements. Prior to the periodic table, it was very difficult to find patterns in the elements. Although many people contributed to the development of the modern Periodic table, it falls mainly down to a handful: Dobereiner Law of Triads:- If you look at the properties and relative atomic masses of 3 elements in group 1:- Element RAM Properties with water Lithium 6.9 Fizzes gently hydrogen given off Sodium 23 Fizzes hydrogen given off Potassium 39.1 Fizzes violently hydrogen given off The RAM of sodium is the mean of lithium and potassium. The rate of reaction of sodium is also in between the other 2. This also occurs for other triads of elements. Newlands Law of Octaves:- Newlands was the first to notice that if you arrange the elements in order of RAM, every 8 th element exhibited similar physical and chemical properties. This only worked for the first 16 elements and other elements were missing. Na Mg Al Si P S Cl Ar K Mendeleev s Periodic Table Mendeleev was the first to put the 60 known elements in the modern arrangement. Mendeleev arranged the elements in increasing mass order but left gaps for elements that had not been discovered yet. Although there were only 60 elements at the time the, gaps he left made it possible to predict the properties of elements that had not yet been discovered. Mendeleev s prediction about Gallium (not yet discovered) was incredibly close:- Mendeleev s prediction Actual properties Atomic mass = 72 Atomic mass = 72.6 Density = 5.5gcm -1 Density = 5.35gcm -1 Light grey colour Silvery grey colour Will combine with 2 atoms of oxygen to form a white powder (the oxide) with a high melting point. Combines with 2 atoms of oxygen to form a white powder (the oxide) with a melting point above 1000 o C Oxide density = 4.7gcm -3 Oxide density = 4.2gcm -3 Will combine with 4 atoms of chlorine to form a chloride with a boiling point less than 100 o C. Will combine with 4 atoms of chlorine to form a chloride with a boiling point less than 100 o C. Chloride density = 1.9gcm -3 Chloride density = 1.8gcm -3 1

2 These predictions were 17 years before Gallium was discovered and shows the correctness of his ideas. Henry Moseley Moseley showed that the real order of the elements in the Periodic Table is based on atomic number and not RAM. The sequence is close but not exactly the same. The modern Periodic Table The elements are arranged in increasing atomic number. Elements with similar properties appear in a regular pattern in the list Groups: A vertical column is called a group These all have the same number of electrons in their outer shells. Groups often show an increase or decrease in similar properties. Periods: A horizontal row is called a period These elements all have the same number of electron shells (if you ignore sub shells) Electron configurations Once the blocks are labelled, any electron structure can be works out. 2

3 Periodicity: Periodicity: The repeating trends in physical and chemical properties of elements as you go across the Periodic Table Periods often show gradual changes in properties. These patterns are repeated across each Period. This is called periodicity Although Periodicity only looks at trends in properties across periods, we will also remind ourselves of trends down groups: Simple examples: Across a Period: Metal à Non metal Across a Period: Electronegativity increases 1) Electron configurations Chemical reactions are due to the outer shell electrons only. This means that those elements with the same number of electrons in the outer shell will react in similar ways. The vigour of the similarities will depend on shielding and the number of protons in the nucleus (Li - Cs). This is what causes Periodicity - these regular repeating patterns: Li Be B C N O F Ne [He]2s 1 [He]2s 2 [He]2s 2 2p 1 [He]2s 2 2p 2 [He]2s 2 2p 3 [He]2s 2 2p 4 [He]2s 2 2p 5 [He]2s 2 2p 6 Na Mg Al Si P S Cl Ar [Ne]3s 1 [Ne]3s 2 [Ne]3s 2 3p 1 [Ne]3s 2 3p 2 [Ne]3s 2 3p 3 [Ne]3s 2 3p 4 [Ne]3s 2 3p 5 [Ne]3s 2 3p 6 2) Atomic radius Across a Period, atomic radius decreases: Shells: Same number of electron shells Shielding: Similar amount of shielding (same number of inner shells) No. Protons: Number of protons increases Attraction: Therefore attraction is greater so shell becomes slightly closer. Down a Group atomic radius increases: Shells: More electron shells Shielding: More shielding (more inner shells) 3

4 Ionic radius Positive ions: These are smaller due to fewer electrons being attracted by the same number of protons attraction increases With Group 2-3 metals, they also lose their outer shell electrons. Negative ions: These are larger due to more electrons being attracted by the same number of protons attraction decreases 3) Ionisation energies These have exactly the same explanation as atomic radius with the addition of Energy required to remove an electron Across a Period Ionisation energy generally increases: Shells: Same number of electron shells Shielding: Similar amount of shielding (same number of inner shells) No. Protons: Number of protons increases Attraction: Therefore attraction is greater Energy: Energy required to remove an electron increases Down a Group Ionisation energy decreases: Shells: More electron shells Shielding: More shielding (more inner shells) No. Protons: Number of protons increases but is outweighed by shells and shielding Attraction: Therefore attraction is less Energy: Energy required to remove an electron decreases Summary for atomic radius and Ionisation energies 4

5 4) Melting and boiling points General trends: Increase in Boiling point from Gp 1-4. Sharp drop from Gp 4-5 Low Boiling points for Gp 5-0 Explanation: As you move from left to right there is a general pattern: Metals à Non metals and Solid à Liquid / Gas There are different structures, forces of attraction and bonding. All 3 of these must be considered. As we move across a period: We shall take each one in turn: Metallic à Giant Molecular à Molecular à Atomic Metallic bonding - Gp 1-3 The boiling points increase as you move along Group 1-3. This is due to: Ø Increase in charge on metal ion Ø The number of delocalized outer electrons Ø Increased attraction so more energy required 5

6 Giant covalent structures - Gp 4 These are carbon and silicon This is due to the strong covalent bonds making up a giant structure. For these to melt / boil: Ø Many strong covalent bonds must be broken Ø Requires lots of energy to do this Carbon - Diamond Carbon - Graphite Silicon Molecular / Atomic - Groups 5-0: These are elements that are simple molecular (N2, O2, P 4, S 8 ) or single atoms (Ne, Ar). For these to melt / boil, Van der Waals must be overcome and these are weak. Van der Waals depend upon: Ø Number of electrons Ø More electrons = greater attraction Ø More energy required to overcome Melting and boiling points of: S 8 > P 4 > Cl 2 S 8 is bigger than P 4 which is bigger than Cl 2 S 8 has more electrons than P 4 which has more electrons than Cl 2 S 8 has stronger VDW than P 4 which has stronger VDW than Cl 2 6

7 Summary: Period 2 Li Be B C N 2 O 2 F 2 He Period 3 Na Mg Al Si P 4 S 8 Cl 2 Ar Atomic radius 1 st Ionisation energy Electronegativity Decreases Increases Increases Structure and bonding Forces Melting / Boiling points Giant metallic Strong electrostatic forces of attraction between positive ions and negative delocalised electrons Giant covalent Strong covalent bonds between atoms Simple molecular / atomic structures weak VDW forces of attraction between molecules / atoms Increases Highest Decreases Questions: 1) State and explain the trend in melting points of the metals Na à Al. Include a diagram. 2) State and explain the trend in boiling points of the elements P 4, S 8 and Cl 2 3) State and explain why the boiling point of Si is the highest across its Period 7