1. 5: Trends in the periodic table j.represent data, in a graphical form, for elements 1 to 36 and use this to explain the meaning of the term ‘periodic property’ k.explain trends in the following properties of the element from periods 2 and 3 of the periodic table: i.melting temperature of the elements based on given data using the structure and the bonding between the atoms or molecules of the element ii.ionization energy based on given data or recall of the shape of the plots of ionization energy versus atomic number using ideas of electronic structure and the way that electron energy levels vary across the period. Connector – Explain what is meant by nuclear charge atomic or ionic radius screening effects ionisation energy electron affinity

2. nuclear charge atomic or ionic radius screening effects The number of protons present in the nucleus of an atom or ion Half the distance between the nuclei of two touching atoms. This is the reduced attraction between the nucleus and the outer energy level electrons when there are electrons in energy levels between them.

3. ionisation energy electron affinity is the amount of energy required to remove a mole of electrons from a mole of gaseous atoms or ions. e.g. Na (g) → Na + (g) + e - is the amount of energy released when a mole of gaseous atoms each gain an electron. e.g.

4. General Periodic Trends We need to consider and explain how the following factors change, either across a period or down a group. Atomic radius Ionic radius Ionisation energy Electron affinity

5. General Periodic Trends – Across a period 1.How does the nuclear charge change across a period? 2.What effect does this have on the electrons in the outer energy level? 3.What is the overall effect on atomic radius across the period?

6. Atomic radius – Across a Period As you move along a period, the nuclear charge becomes increasingly positive as the number of protons in the nucleus increase. Although the number of electrons also increases, the outer electrons are all in the same energy level. This means that electrons are attracted more strongly to the nucleus, thus reducing the atomic radius across a period.

7. General Periodic Trends – Down a group 1.How does the arrangement of the electrons in an atom change going down a group? 2.What happens to the nuclear charge? 3.What are the effects of 1 & 2 on the electrons in the outer energy level? 4.What is the overall effect on atomic radius down a group?

8. As you go down a group, the outer electron(s) enter into a new energy level. Although, the nucleus also gains positive protons(charge), the electrons are both further away and screened by more electron energy level. As a result, electrons are not held so tightly, and the atomic radii increases. Atomic radius – Down a Group

9. Atomic Radius: the radius of an atom in picometers 1 2 13 14 15 16 17 18

10. Atomic radius vs. atomic number

11. Which is Bigger? Na or K ?Na or K ? Na or Mg ?Na or Mg ? Al or S ?Al or S ?

12. Does the size go up or down when losing an electron to form a cation? Does the size go up or down when losing an electron to form a cation? Ionic radius

13. CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Ionic radius Forming a cation.

14. Does the size go up or down when gaining an electron to form an anion? Does the size go up or down when gaining an electron to form an anion? Ionic radius

15. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes.Trends in ion sizes are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p - Ionic radius

16. Trends in Ion Sizes

17. Which is Bigger? Cl or Cl - ?Cl or Cl - ? K + or K ?K + or K ? K + or Ca +2 ?K + or Ca +2 ? I - or Br - ?I - or Br - ?

18. Periodicity The repeating patterns or trends are known as the periodicity of the elements and the properties are known as periodic properties.

19. Ionization Energy It is the amount of energy required to remove a mole of electrons from a mole of gaseous atoms or ions.

20. Ionization Energy Ionisation energy is affected by three main factors: The attraction between the nucleus and the outermost electron The size of the positive nuclear charge Screening effect

21. Ionization energy vs. atomic number

22. Which has a higher 1 st ionization energy? Mg or Ca ? Al or S ? Cs or Ba ?

23. On the graph, the end of each period is marked by a peak of a high I.E. of a noble gas This distinctively high I.E., as we have seen, is due to the fact that noble gases have a complete outer shell (the most stable electronic configuration

24. This graph shows that the first I.E. doesn’t increase smoothly. This is due to the presence of subshells within shells. In the period Li to Ne, Be has a higher I.E. than B. Same when comparing Mg to Al. This is because, removing an electron from B or Al, would mean taking it from a 2p subshell. However, taking one from Be or Mg, means removing from a full s subshell. These are stable, so more energy required.

25. Also, N and P also have peaks in the graph. This is because, they have half filled p subshells. Again underlying how the most stable configuration is when you have a full or half-full shell or subshell. Its graphs like this that prove the existence of subshells.

26. ELECTRON AFFINITY A few elements GAIN electrons to form anions. 1) Electron affinity is the energy change which occurs when an electron is accepted by an atom in the gaseous state. A (g) + e  A - (g)

27. a) Ionization Energy is always endothermic. It always takes energy to remove an electron. Change in Enthalpy (∆ H) is always +ve. b) Electron Affinity can be either endothermic or exothermic depending on the element. c) An exothermic (Change in Enthalpy is -ve) value for the electron affinity indicates that energy is released upon the addition of an electron to a gaseous atom. d) The greater the negative value of the electron affinity, the greater the tendency of an atom to accept an electron.

28. e) A +∆H indicates that energy must be absorbed for an atom to gain an electron. f) As we go from left to right on the periodic table, the elements have, in general, an increasing tendency to form negative ions.

29. Electron Affinity of Oxygen ∆E is EXOthermic because O has an affinity for an e-. [He]      O atom EA = - 141 kJ + electron O [He]       - ion

30. Electron Affinity of Nitrogen ∆E is zero for N - due to electron- electron repulsions. EA = 0 kJ [He]     Natom  [He]    N - ion  + electron

31. Spot the trend of Electron affinity in a group and in a period

32. Electron Affinity Trends (Same as for Ionization Energy) Group Trends: Electron affinity decreases from top to bottom on the periodic table Outermost electrons are further away from the nucleus and therefore easier to remove Shielding effect increases down the group Period Trends: increases from left to right Nuclear charge is increasing with no increase in shielding effect Outermost electrons are closer to the nucleus

33. Trends in Electron Affinity

34. Electron affinity MORE STABLE LESS STABLEMORE STABLE

35. Electron affinity

36. Melting Temperature Melting Temp. (for a substance) = tamp at which pure solid is in equilibrium with pure liquid at atmospheric pressure. This is determined by the packing and bonding of atoms in a substance.

37. Melting Temperature Metals like Li, Mg, and Al have peaks due to strong metallic bonding (ions packed in sea of electrons) Si and C are giant covalent structures with extremely strong covalent bonding. So at peaks. After peaks, the elements tend to exist as simple covalent molecules (diatomic etc). Therefore, despite strong intramolecular forces (covalent bond) have weak intermolecular attraction. So easy to pull individual molecules apart.

38. Electronegativity Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The trend across Period 3 looks like this: Explaining the trend As you go across the period, the bonding electrons are always in the same energy level. They are always being screened by the same inner electrons. All that differs is the number of protons in the nucleus. As you go from sodium to chlorine, the number of protons steadily increases and so attracts the bonding pair more closely.

39. Homework Homework task: Due date: Criteria for Grade C: Criteria for Grade B: Criteria for Grade A/A*: