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Section 2

Section 2. Periodicity. Bonding in the Elements 1-20 (a). L.I. To learn about Bonding in the Elements 1-20 S.C. By the end of this lesson you should be able to describe the metallic bond explain what is meant by the term monatomic explain what London dispersion forces are and how

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Section 2

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  1. Section 2 Periodicity Bonding in the Elements 1-20 (a)

  2. L.I. To learn about Bonding in the Elements 1-20 • S.C. By the end of this lesson you should be able to • describe the metallic bond • explain what is meant by the term monatomic • explain what London dispersion forces are and how • they arise • explain what happens to the strength of LDF as the • atom size increases • explain the difference between covalent network and • covalent molecular in terms of bpt and mpt • give examples of metallic, covalent molecular, covalent • network and monatomic elements

  3. Periodic Pattern • Johan Wolfgang Dobereiner – triads the atomic mass of the central element was approximately the mean of the other two.

  4. What does the periodic table and the sound of music have in common? • John Newlands – octaves based on atomic mass (musical notes). o • Every eighth element showed similarities

  5. The modern Periodic Table is based on the work of Dimtri Mendeleev in 1869 • He arranged the elements based on: atomic mass, similar properties • He left gaps and made predictions for missing elements fun activity in workbook

  6. Bonding in metals activity in workbook • Metallic bonding is the electrostatic attraction between the positively charged ions and the delocalised electrons.

  7. Bonding in metals The outer electrons are delocalised and free to move throughout the lattice, making metals good conductors of electricity. The greater the number of electrons in the outer shell the stronger the metallic bond. So the melting point of Al>Mg>Na

  8. ++ Bonding in Monatomic elements He Noble gases have full outer electron shells They do not need to combine with other atoms. The noble gases occur as single atoms, they are said to be monatomic. Since they can be liquefied and solidified there must be some weak attraction between the atoms.

  9. δ+ δ+ δ+ δ- δ- δ- ++ ++ Bonding in monatomic elements The electrons in an atom “wobble” and become unevenly distributed causing one side of the atom becomes slightly negative while the other side becomes slightly positive. These slight charges are given the symbol δ ‘delta’. A temporary dipole is therefore formed. A dipole can induce other atoms to form dipoles, resulting in weak attractions between particles. activity in workbook London dispersion forces

  10. Bonding in monatomic elements London dispersion forces are a type of van der Waal force. They are very weak attractive forces.

  11. Comparing Boiling points The melting and boiling point of a substance gives an indication of the strength of the forces of attraction holding atoms or molecules together. activity in workbook

  12. 180 166 160 140 120 100 80 60 40 20 0 121 4 87 27 Noble gases b.p.’s Helium Neon b.p / K Argon Krypton Xeon activity in workbook B.p.’s increase as the size of the atom increases This happens because the London dispersion forces increases with increasing size of atoms.

  13. Covalent Molecular Elements Many non-metals exist as discrete covalent molecules held together by covalent bonds. Discrete molecules have a definite number of atoms bonded together. Fluorine atom Fluorine molecule F2 diatomic 9+ 9+ 9+ activity in workbook

  14. Examples of discrete molecules: Cl Cl activity in workbook - table

  15. weak London dispersion forces strong covalent bonds Melting point low – why? activity in workbook - paragraph

  16. Comparing Boiling points activity in workbook

  17. 500 450 400 350 300 250 200 150 100 50 0 457 332 85 238 Halogens b.p.’s Fluorine b.p./ K Chlorine Bromine Iodine activity in workbook As the size of the halogen molecule increases the boiling point increases. The bigger the molecule the stronger the London dispersion forces between the halogen molecules.

  18. Fullerenes, molecules of carbon Fullerenes exists as large covalent molecules with a definite formula. Fullerenes were discovered in 1985 by Buckminster Fuller. Fullerenes are spherical in shape and usually contain sixty or seventy carbons. C60 is known as Buckminster fullerene activity in workbook

  19. Covalent Network Elements Carbon - diamond Diamond has a covalent network structure Each of the outer electrons in a carbon atom can form a covalent bond with another carbon atom. So every C bonds to 4 others. m.p.’s C > 3642oC It is high because many covalent bonds have to be broken.

  20. Carbon - Graphite Carbon bonded to only 3 other Carbons The spare (4th) electron is delocalised and so free to move. Graphite is a conductor of electricity. Van der Waals forces between the layers allows layers to slide over each other. Graphite can be used as a lubricant

  21. Properties of graphite and diamond

  22. Other Network Structures In the first 20 elements, only Boron, Carbon and Silicon have covalent network structures. m.p.’s B 2300oC, C > 3642oC and Si 1410oC activity in workbook

  23. BONDING IN ELEMENTS - A SUMMARY activity in workbook

  24. B C Si Bonding patterns of the 1st 20 elements H He Covalent Molecular Metallic lattice B N N C O O F F Ne Li Li Be Be Monatomic P Si S Cl Ar P S Cl Na Na Mg Mg Al Covalent Network K K Ca Ca C , in the form of fullerenes, is covalent molecular

  25. http://www.ltscotland.org.uk/highersciences/chemistry/ animations/bonding_structure.asp This interactive animation provides a visual representation of the bonding and structure of the first twenty elements in the periodic table, taking into account both the intra- and inter-molecular forces involved.

  26. Questions on elements – bonding and structure • Explain why the covalent network elements have high melting and boiling points. • Explain why the discrete molecular and monatomic elements have low melting and boiling points. • Does diamond conduct electricity? Explain. • Does graphite conduct electricity? Explain. • How does the hardness of diamond compare with graphite? Explain. • Give a use for both diamond and graphite. • Complete the following table:

  27. Questions on elements – bonding and structure 7. Complete the following table:

  28. Section 2 Periodicity Patterns in the Periodic Table (b)

  29. L.I. To learn about covalent radius • S.C. By the end of this lesson you should be able to • describe the term covalent radius • explain the changes in covalent radius down a group • explain the changes in covalent radius across a period • explain why there is no stated covalent radius for the • noble gases

  30. Covalent Radius The size of an atom is indicated by its covalent radius. (Page 7 of data booklet). There is no definite edge to an atom. 266pm However, bond lengths can be worked out. The covalent radius of an element is half the distance between the nuclei of 2 of its bonded atoms From above the covalent radius would be 133 pm. Covalent radius – picometres (pm) 1pm =1 X 10 – 12 m

  31. Trends in covalent radius - Across a period activity in workbook Na 154 pm, Mg 145 pm, Al 130 pm, Si 117 pm, P 110 pm, S 102 pm Going across a period the covalent radius (atomic size) decreases. Why? Going across a period the nuclear charge increases. The attraction between the outer electrons and the positive nucleus increases. Thus the outer electrons are more strongly attracted and so the atom size is smaller.

  32. Li 134 pm Na 154 pm K 196 pm Rb 216 pm Trends in covalent radius – Down a group activity in workbook Going down a group the covalent radius (atomic size) increases. On moving down a group from one element to the next the number of electron shellsincreases. So the outer electrons are further from the nucleus and the atom size increases.

  33. Why is there no covalent radius value for the noble gases?

  34. L.I. To learn about ionisation energies • S.C. By the end of this lesson you should be able to • describe the term 1st ionisation energy • write equations for the 1st ionisation energy • explain the trend in 1st ionisation energy down a group • explain the trend in 1st ionisation energy across a period • describe the term 2nd ionisation energy • carry out calculations involving ionisation energy

  35. Ionisation energies The first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. Units are kJmol-1. This is an endothermic process. (Page 11 of the data booklet.) Na(g) Na+ (g) + e

  36. Cl(g) Cl+ (g) + e

  37. Trends in 1st ionisation energy – Across a period activity in workbook Going across a period the ionisation energy increases. Going across a period the nuclear chargeincreases. The attraction between the negative electrons and the positive nucleus increases. Thus the electrons are more tightly held and so more energy is needed to remove the outer electrons.

  38. Trends in 1st ionisation energy – Down a group activity in workbook Going down a group the ionisation energy decreases. The explanation for this is (i) on moving down a group from one element to the next the number of electron shellsincreases and so the outer electron is further from the nucleus and less tightly held. (ii) the inner shells provide a screening effect which also decreases the attractive forces between the outer electrons and nucleus.

  39. The 2nd Ionisation Energy The second ionisation energy of an element is the energy required to remove the second mole of electrons. Second Ionisation First Ionisation Mg(g)  Mg+(g) + e- Mg+(g)  Mg2+ + e- ΔH = +738 kJ mol-1 ΔH = +1451 kJ mol-1 Third Ionisation Mg2+(g)  Mg3+ + e- ΔH = +7733 kJ mol-1 activity in workbook

  40. L.I. To learn about electronegativity • S.C. By the end of this lesson you should be able to • describe the term electronegativity • explain the trend in electronegativity down a group • explain the trend in electronegativity across a period • explain why there are no quote values of electronegativity • for the noble gases

  41. Electronegativity The electronegativity is a measure of the attraction an atom involved in a bond has for the shared pair of electrons. Electronegativity values are based on the Pauling Scale, devised by Linus Pauling an American Chemist. Values on the Pauling Scale range from 0 to 4. A list of these values can be found in the data booklet on page 11. The higher the number on the Pauling scale is, the greater the attraction an atom has for the bonding electrons.

  42. Electronegativity values can be useful in predicting which type of bonding is most likely between two elements. (More about this later) Electronegativity – Across a Period activity in workbook On crossing a period, electronegativity values increase. This is caused by an increase in nuclear charge as you move across a period from left to right. Electronegativity – Down a Group activity in workbook As you go down a group, electronegativity values decrease. This is caused by the addition of another energy level of electrons as you go down a group which shields the bonded electrons from the nucleus; therefore they are not attracted as strongly.

  43. Electronegativity - The Monatomic Gases Why no values for group 8 elements?

  44. Section 3 Structure and Bonding Bonding in Compounds

  45. L.I. To learn about bonding in compounds (a) • S.C. By the end of this lesson you should be able to • describe the bonding and structure in ionic compounds • explain the melting point of ionic compounds • describe the bonding and structure in covalent network • compounds • explain the melting point of covalent network compounds • describe the bonding and structure in covalent molecular • compounds • explain the melting point of covalent molecular compounds

  46. Three different types of compound - ionic,covalent molecular or covalent network. Ionic Bonding Na Cl Na+ + Cl- 2)8)1 2)8)7 2)8 2)8)8 In ionic compounds atoms achieve a full outer shell by either losing or gaining electrons and so form charged particles called ions.

  47. Complete for sodium, chlorine, bromine, oxygen, aluminium and nitrogen. Metal atoms always lose electrons to form positive ions e.g Na+ Non-metal atoms always gain electrons to form negative ions e.g F- Glow: ionic bonding ionic compounds

  48. On show me boards – work out how these elements form an ionic compound Sodium chloride Lithium fluoride Magnesium oxide Aluminium nitride Calcium chloride Now write ionic formula for the above.

  49. Ionic Bonding + - ------- ionic bond NaCl 3D lattice – regular repeating pattern of ions The attraction between positive and negative ions holds the compound together. The electrostatic attraction between positive and negative ions is an ionic bond.

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