1. CHAPTER 10: PERIODICITY

2.  In the periodic table elements are arranged in order of ……  The elements are in a certain group and period  Elements in a group (example 1) have similar properties  Elements in a period have periodicity The table is split into blocks; in each block the elements are filling, or have just filled, particular types of orbital s block1 and 2 end in s 1 or s 2 p block13, 14, 15, 16, 17 and 18 end in p 1 to p 6 d blockTransition elements (3-12) end in d 1 to d 10 f blockActinides and Lanthanides end in f

3. Because many physical and chemical properties are influenced by the outer shell configuration of an atom, we see these properties across a period: 1. atomic radius 2. ionic radius 3. melting point and electrical conductivity 4. ionisation energy 5. electronegativity 6. electron affinity 7. boiling point Periodicity and physical properties

4. Refresh your memory: ELECTRONIC CONFIGURATION

5. The Aufbau principle states that… “ELECTRONS ENTER THE LOWEST AVAILABLE ENERGY LEVEL” In period 3 the electrons fill the 3s orbital first, followed by the 3p orbitals. Notice how the electrons in the 3p orbitals remain unpaired, if possible, according to Hund’s Rule.

6. Moving from left to right the elements go from metals through metalloids with giant structures to the simple molecular structure of non-metals. NaMg AlSiP 4 S 8 Cl 2 Ar metalloid Typical properties Metals Non-metals Appearance solids - shiny when cut gases, liquids, dull solids Hardness malleable and ductile brittle Electrical conductivityexcellent poor Melting point high low Period 3 elements

7. “We can compare the size of different atoms using their atomic radii.” The data for these measurements can come from an element’s single covalent radius. (There are other ways of measuring the atomic radius but because most elements in the periodic table combine using covalent bonds here this method is applied). This is not actually measuring the true radius of an atom. In metals you measure metallic radius (half the distance between the inter- nuclear distance of what are effectively ions). Covalent radius is half the distance between the nuclei of atoms joined by a covalent bond. 1. Atomic radius

8. We study the following properties across a period: 1. Atomic radius 2. Ionic radius 3. Melting point 4. Electrical conductivity 5. Ionisation energy 6. Electronegativity 7. Electron affinity Periodicity 1.Reactions with chlorine gas 2.Oxidation numbers 3.Acidic/basic behavior

9. Decreases across a given period The nuclear charge increases by +1 each time. As the nuclear charge increases it has a greater attraction for the electrons which, importantly, are going into the same principal quantum shell and pulls them in slightly. Trend in atomic radius

10.  We know that metals will loose electrons if they become ions hence the atomic radius of an ion of a metal is …… than its original atom.  We also know that non-metals will gain electrons when becoming an ion, hence their ionic radii will be ….. than their original atom. 2. Trend in ionic radius

11. Time for some practice Please make Q. 1 and 2

12. ELECTRICAL CONDUCTIVITY

13. Na Mg Al Si P S Cl Ar 0.4 0 Substances conduct electricity when ions or electrons are free to move. Periods Overall decrease across periods Na, Mg, Almetallic bonding with delocalised electrons Si, P, S, Clcovalently bonded - no electrons are free to move Ar monatomic - electrons are held very tightly UNITS:- Siemens per metre 3. Electrical conductivity

14. MELTING POINT

15. Na Mg Al Si P S Cl Ar 3000 0 Boiling and melting points are a measure of the energy required to separate the particles in a substance. Bond type is significant. Metals Na-Al Melting point increases due to the increasing strength of metallic bonding caused by the larger number of electrons contributing to the “cloud” Kelvin The electron cloud in magnesium is denser than in sodium so more energy is required to separate the ‘ions’ SODIUM MAGNESIUM

16. Na Mg Al Si P S Cl Ar 3000 0 Metalloid SILICON Large increase in melting point as it has a giant molecular structure like diamond A lot of energy is required to break the many covalent bonds holding the atoms together. Kelvin Melting point

17. Na Mg Al Si P S Cl Ar 3000 0 P, S, Cl are simple covalent molecules Melting point depends on the weak intermolecular van der Waals’ forces. The larger the molecule the greater the van der Waals’ forces Kelvin Melting points

18. Na Mg Al Si P S Cl Ar 3000 0 PHOSPHORUS can exist is several allotropic forms. In red phosphorus, each molecule exists in a tetrahedral structure. The atoms are joined by covalent bonds within the molecule formulaP 4 relative mass124 melting point44°C Melting point drops dramatically as intermolecular attractions are now due to weak van der Waals’ forces. Kelvin Melting points

19. Na Mg Al Si P S Cl Ar 3000 0 SULPHUR can exist is several allotropic forms. Molecule can exist in a puckered eight membered ring structure. The atoms are joined by covalent bonds within the molecule formulaS 8 relative mass256 melting point 119°C Melting point rises slightly as the molecule is bigger so has slightly stronger van der Waals’ forces. Kelvin Melting points

20. Na Mg Al Si P S Cl Ar 3000 0 CHLORINE Exists as a linear diatomic molecule. The atoms are joined by covalent bonds within the molecule formulaCl 2 relative mass71 melting point -101°C Melting point falls slightly as the molecule is smaller so has slightly lower van der Waals’ forces. Kelvin Melting points

21. Na Mg Al Si P S Cl Ar 3000 0 ARGON Exists as a monatomic species. formulaAr relative mass40 melting point -189 °C Melting point falls. Kelvin Melting points

22. Na Mg Al Si P S Cl Ar 3000 0 Boiling points tend to be a better measure and show better trends because solids can be affected by the crystal structure as well as the type of bonding. As is expected, the boiling points are higher than the melting points. Kelvin Boiling points across period 3

23. 1st IONISATION ENERGY

24. 1st Ionisation Energy INCREASES across a period Nuclear charge increases across the period. Each extra electron is going into the same shell. Electrons are held more strongly and are harder to remove. However 2 exceptions. The energy required to remove ONE MOLE of electrons (to infinity) from ONE MOLE of gaseous atoms to form ONE MOLE of gaseous positive ions. e.g.Na(g) Na + (g) + e - Al(g) Al + (g) + e - Make sure you write in the (g) 1 st IE

25. INCREASES across a period Nuclear charge increases by one each time. Each extra electron, however, is going into the same main energy level so is subject to similar shielding and is a similar distance away from the nucleus. Electrons are held more strongly and are harder to remove. However the trend is not consistent. 3s 3p Na11+ Mg12+ Al13+ Si14+ P15+ S16+ Cl17+ Ar18+ NUCLEAR CHARGE Trend in 1 st IE across period 3

26. Na Mg Al Si P S Cl Ar 1500 1000 500 3s 3p There is a drop in the value for sulfur. The extra electron has paired up with one of the electrons already in one of the 3p orbitals. The repulsive force between the electrons means that less energy is required to remove one of them. There is a drop in the value for aluminium because the extra electron has gone into a 3p orbital. The increased shielding makes the electron easier to remove. Trend in 1 st IE

27. Time for some practice Please make Q. 3 and 4 on page 153

28. ELECTRONEGATIVITY

29. Electronegativity scale

30. Na Mg Al Si P S Cl Ar 3.5 3.0 2.5 2.0 1.5 1.0 0.5 A measure of the attraction an atom has for the pair of electrons in a covalent bond. Increases across a period... because the nuclear charge is increasing and therefore so does the attraction for the shared pair of electrons in a covalent bond. UNITS:- Pauling Scale “The ability of an atom to attract the pair of electrons in a covalent bond to itself.”

31. CHEMICAL PROPERTIES OF PERIOD 3 ELEMENTS

32. Oxygen has a higher electronegativity than all the elements of period 3, hence the elements in the oxides of period 3 elements exist in positive oxidation states Element: NaMgAlSiPSClAr Oxide: Na 2 O Oxidation +1 No. MgOAl 2 O 3 SiO 2 P 4 O 10 SO 2, SO 3 Cl 2 O, Cl 2 O 7 +2+3+4+5+4, +6+2, +7 Refresh: oxidation states

33. Refresh more Na 2 O MgO Al 2 O 3 SiO 2 P 4 O 10 SO 2, SO 3 Chemical ionic Bonding Structure giant ionic Conduct in good Liquid state ionic giant ionic good ionic covalent giant ionic good covalent giant covalent none covalent simple molecular none covalent simple molecular none

34. 2Na(s) + Cl 2 (g) 2NaCl(s)(vigorous reaction) Mg(s) + Cl 2 (g) MgCl 2 (s)(vigorous reaction) 2Al(s) + Cl 2 (g) Al 2 Cl 6 (s)(vigorous reaction) Si(s) + Cl 2 (g) SiCl 4 (l)(slow reaction) 2P(s) + 5Cl 2 (g) 2PCl 5 (l)(slow reaction) Sulfur can exist as SCl 2 and S 2 Cl 2. Argon does not react with Chlorine Trend in reactions We study the reaction of elements with Cl 2 because due to its high reactivity, it is commonly found in nature bonded to many different elements

35. Time for some practice Please make Q. 5 on page 155

36. REACTIONS OF OXIDES WITH WATER

37. Na 2 O(s) + H 2 O(l) 2NaOH(aq)(strong alkaline) MgO(s) + H 2 O(l) Mg(OH) 2 (s)(weak alkaline) Al 2 O 3 (s) + H 2 O(l) (no reaction) SiO 2 (s) + H 2 O(l)(no reaction) P 4 O 10 (s) + 6H 2 O(l) 4H 3 PO 4 (aq)(acidic solution) SO 2 (g) + H 2 O(l)H 2 SO 3 (aq)(acidic solution) SO 3 (g) + H 2 O(l)H 2 SO 4 (aq)(acidic solution) Period 3 element oxides with water

38. Al 2 O 3 (s) + 3H 2 SO 4 (aq)Al 2 (SO 4 ) 3 (aq) + 3H 2 O(l) Al 2 O 3 (s) + 2NaOH(aq) + 3H 2 O2NaAl(OH) 4 (aq) Aluminium oxide can react with an acid as well as a base, so it can behave like a base or acid. We call this kind of species that can act as both acids and bases amphoteric. SiO 2 (s) + 2NaOH(aq) Na 2 SiO 3 (aq) + H 2 O(l) Silicon dioxide is insoluble in water, water cannot break down its giant molecular structure. However it acts as an acid in hot concentrated alkali. Aluminum oxide and Silicon dioxide

39. Time for some practice Please make Q. 6 and 7 on page 158/160