2. The Periodic Table

3. Overview

4. Lesson 1: The Periodic Table  Objectives: Reflect on prior knowledge of the periodic table Understand the structure and purpose of the periodic table

5. Reflecting on the Periodic Table  What is the periodic table and what is supposed to show?

6. Dmitri Mendeleev’s Periodic Table The one that started it all off.

7. Wide Format Periodic Table Shows true position of the f-block (lanthanides and actinides)

8. Janet Periodic Table Elements arranged in order of orbital filling. Used frequently by physicists.

9. Benfey Periodic Table Spiral form shows the steady increase in atomic number.

10. Stowe Periodic Table Emphasises the symmetrical nature of the increase in quantum numbers.

11. Zymaczynski Periodic Table Another way to show the symmetry in the underlying quantum numbers.

12. Giguere Periodic Table A 3D representation emphasising the s, p, d and f blocks

13. “Mayan” Periodic Table

14. The Traditional Based on Mendeleev’s work. Easiest to use and display.

15. The Structure of the Periodic Table PERIODS GROUPS

16. Groups and Periods  Groups Elements show similar chemical properties Elements show similar trends in their chemical properties  Periods As you move across periods, changes in the chemical and physical properties that are repeated in the next period This is what ‘period’ and ‘periodic’ refers to

17. Before we get Started  Ionization energy: the amount of energy it takes to strip away the first valence electron  Electronegativity: a measure of how tightly an atom holds onto its valence electrons  Nuclear charge: the attractive force between the positive protons in the nucleus and the negative electrons in the energy levels. The more protons, the greater the nuclear charge.  Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the valence electrons and the nucleus, the more shielding.

18. The periodic table and electron configuration  How does an element’s position in the PT relate to its electron configuration?

19. Key Points  The periodic table arranges the elements according to: Their chemical properties Their electronic structure

20. We Are Here

21. Physical Properties  Objectives: Identify and explain the trends in the physical properties of the first 20 elements including: ○ Atomic radius ○ Ionic radius ○ First ionisation energy ○ Electronegativity ○ Melting point

22. Atomic Radius  This is the ‘size’ of an atom  There is no simple measure as atoms do not have a well defined ‘edge’  We use the: covalent radius This is half the distance between the nuclei of two atoms in a covalent bond This means we don’t have values for the noble gases as they do not form bonds  Values range from are measured in picometres ( 1 pm = 1x10 -12 metres…a thousand-billionth of a metre) and range over: 270 picometres Francium 30 picometres for Hydrogen (helium would be smaller but does not form covalent bonds to be measured)  The main factors influencing atomic radius are: Number of shells (the principal quantum number) The charge in the nucleus

23. Ionic Radius  This is the ‘size’ of an ion and is measured in a similar way to atomic radius  It is measured in a picometres with values ranging over: 272 pm for the Ge 4- ion 16 pm for the B 3+ ion  The main factors influencing ionic radius are: Number of shells (the principal quantum number)…don’t forget this can be affected by the type of ion formed The charge in the nucleus

24. First ionisation energy  This is the energy required to remove one mole of electrons from one mole of gaseous atoms to form positive ions i.e.: A(g)  A + (g) + e -  Values range over: 393 kJ mol-1 for Caesium 1681 kJ mol-1 for Helium  Values are positive because this is an endothermic process  Values are influenced by: Number of inner electron shells (and their shielding) Charge on the nucleus HL: At the finest level – repulsion between electrons in their orbitals

25. Electronegativity  This is a measure of the degree to which an element attracts the shared pair of electrons in a covalent bond Again, this means there are no values for the noble gases  Values range over: 4.0 for Fluorine 0.7 for Francium  Values are influenced by: Number of inner electron shells (and their shielding) Charge on the nucleus  Values are unit-less as this is a relative measure

26. Melting Point  This is the temperature (in Kelvin…i.e. Celsius + 273) at which an element melts  Values range over: 3935 K for Carbon 1 K for Helium  Values are influenced by: Nature of bonding: giant covalent, giant ionic, metallic Strength of bonding Strength of intermolecular forces

27. Reactivity of Chemicals  METALS In METALS reactivity increases as you go down a group because the farther down a group of metals you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity. ACROSS a Period: In METALS reactivity decreases as you go across a period because though they still want to give away valence electrons they have more of them to get rid of, which requires more energy.

28. Reactivity of Chemicals  NON-METALS Reactivity INCREASES as you go UP a Group because the higher up and to the right atoms are, the higher the electronegativity, resulting in a more vigorous exchange of electrons. ○ Fluorine? A greedy, impatient beast when it comes to electron exchange manners. Reactivity INCREASES as you go ACROSS a Period because (notice how trends repeat?) the closer you get to fulling your s- and p- orbitals the more motivated you are to do so.

29. In Summary

30. We Are Here

31. Chemical Trends  Members of a group often have very similar reactivity.  You probably know that carbon will react with hydrogen to form methane, CH 4  You probably did not know that silicon will also react with hydrogen to form silane, SiH 4  Watch this demonstration to see some silane being made https://www.youtube.com/watch?v=qUvNU7EIqMM

32. Three reactions to know  The Group I (alkali) metals react with water as follows: Metal + Water  Metal Hydroxide + Hydrogen  The Group I (alkali) metals react with halogens (Group VII) as follows: Metal + Halogen  Metal Halide  Halogens can react with halide ions as follows (using the example of bromide and chlorine): Bromide + Chlorine  Chloride + Bromine