CfE Higher Chemistry

1. CfE Higher Chemistry Unit 1 Chemical Changes and Structure LI: Describe Trends in The Periodic Table

2. Atomic Structure Revision • Atoms are made up of 3 different types of sub-atomic particles. • These particles are called protons, neutrons and electrons. • An atom basically has 2 areas where these particles exist – the nucleus and the energy levels. • The nucleus has the same spelling as the nucleus of a cell but a different purpose. Nucleus just means centre.

3. The Structure of the Atom How big is an atom?

4. Electron Arrangement Cl chlorine has 17 electrons 17 • First level can take 2 electrons • Second level can take 8 electrons • When the level is half full the electrons have to pair up 17+ • Third level can take 8 electrons • Again half fill with single electrons • The last 3 electrons have to pair up • Chlorine has 1 unpaired electron. This gives an electron arrangement of 2,8,7

5. The Modern Periodic Table • Dmitri Mendeleev arranged the known elements in order of increasing atomic mass • He also arranged the elements in columns depending on their chemical properties e.g. alkali metals and halogens • He left gaps for elements yet to be discovered • He predicted the properties of these elements e.g. germanium Dimitri Mendeleev

6. The Modern Periodic Table We now know that • The atomic number tells us the number of protons in the nucleus of each element • The elements in each column have the same number of outer electrons • The number of outer electrons has an effect on the reactivity of the elements Group 1 Reactivity of Group 1 – the Alkali metals

7. Trends in the Periodic Table • There are trends in the properties of the elements going • Across a period • Down a group

8. Atomic Size Learn this!! • Atomic size is measured as covalent radii • This is the prediction of the size of a single atom • It is half the distance between covalently bonded atoms • Using the data book, plot atomic size (vertical axis) against atomic number (horizontal axis) for elements 3 to 20 • Use a dotted line between elements 9 and 11, and 17 and 19

9. Start of a new row Covalent radius increases down a group Covalent radius decreases across a period

10. Covalent Radii • How does it change across a period? • Decreases • How does it change down a group? • Increases • Atomic size is a periodic property • It follows a pattern going across and down the Periodic Table • Why are there no values for the noble gases? • As they do not form covalent bonds, they cannot have a covalent radius

11. Decreasing Covalent radius

12. Covalent Radius - Across a Period • Why does atomic size decrease going across a period? • Think about what happens to the atoms as we move across? • What changes in the nucleus? • What changes in the energy level?

13. Atomic Size 3 protons in nucleus 9 protons in nucleus Lithium Fluorine

14. Atomic Size electrons 3 protons in nucleus 9 protons in nucleus Lithium Fluorine

15. Atomic Size Learn this!! • Nuclear charge increases as we go across a period – we add protons to the nucleus • The added electrons join the same electron shell, each electron is subject to an increasing nuclear attraction • The greater the nuclear charge • the greater the attraction of the electrons • The closer the electrons are drawn to the nucleus • The smaller the atomic size

16. Atomic Size • Why does atomic size increase going down a group? • What happens to • The nucleus • The electrons?

17. 11p Atomic Size 3p 11p sodium lithium

18. Atomic Size Learn this!! • As we move down a group the atomic size increases • This is because another level of electrons is added • The filled inner energy levels shield the outer electrons from the pull of the nucleus

19. Decreasing covalent radius The nuclear charge increases and electrons are added to the same energy level. Energy levels are pulled closer to the nucleus so covalent radius decreases. Increasing covalent radius The number of energy levels increases. The nuclear charge increases but the outer energy levels are shielded from the effect of the nucleus by the inner energy levels.

20. Ionisation Energy Learn this!! • Ionisation energy is the energy required to remove one electron from each atom in a mole of gaseous atoms of the element E(g) → E+(g) + e- where E is any element

21. First Ionisation Energy “The energy required to remove one electron from each atom in one mole of free gaseous atoms” Mg(g) → Mg+(g) + e Exothermic or endothermic? H1 = +738 kJmol-1

22. Ionisation Energy • Using the data booklet • Draw a graph of first ionisation energy against atomic number for the first 20 elements • Use a dotted line between the noble gases and the elements in group 1

23. Start of a new row Ionisation energy increases across a period Ionisation energy decreases down a group

24. Ionisation Energy – Across a Period Learn this!! • How does it change going across a period? • The ionisation energy increases as we move across, due to • The increased nuclear charge • The decrease in atomic size • Electrons are held more tightly • More energy required to remove one electron

25. Ionisation Energy – Down a Group Learn this!! • How does it change going down a group? • The ionisation energy decreases as we move down a group, due to • The shielding (or screening) effect of each added energy level • The increased distance from the nucleus • Electrons are held less tightly • Easier to remove one electron

26. Increasing ionisation energy The nuclear charge increases and electrons are added to the same energy level. Energy levels are pulled closer to the nucleus. Decreasing ionisation energy The number of energy levels increases. The nuclear charge increases but the outer energy levels are shielded from the effect of the nucleus by the inner energy levels.

27. Ionisation Energy • Where in the Periodic Table are we most likely to find atoms that form positive ions? • The bottom left hand side (easiest to remove an electron) • Where is it hardest to remove an electron? • The top right hand side • A full electron shell, close to the nucleus • i.e. the noble gases

28. Ionisation Energies • The ionisation energies listed in the data booklet are all for the removal of one electron • More than one electron can be removed from an atom and so we have second, third and fourth ionisation energies

29. Ionisation Energies First: “The energy required to remove one electrons from each atom in one mole of free gaseous atoms” Mg(g) → Mg+(g) + e H1 = +738 kJ mol-1 Second: “The energy required to remove a second electron from each 1+ ion in one mole of free gaseous atoms” Mg+(g) → Mg2+(g) + e H2 = +1450 kJ mol-1 Third: “The energy required to remove a third electron from each 2+ ion in one mole of free gaseous atoms” Mg2+(g) → Mg3+(g) + e ∆H3 = + 7750 kJ mol-1

30. Questions on Ionisation Energy Explain the large difference between the 1st and 2nd ionisation energies for Group 1 metals Na 1st 502 kJmol-1 2nd 4560 kJmol-1 K 1st 425 kJmol-1 2nd 3060 kJmol-1 Look at the electron arrangement of the atoms Na 2)8)1 K 2)8)8)1 1 electron in outer shell – relatively easy to remove Atoms now have full outer shell – stable – far more difficult to remove further electrons

31. Questions on Ionisation Energy The same thing happens with the other groups Once the outer electrons are removed there is a big jump in ionisation energies to remove electrons from full levels closer to the nucleus

32. Questions on Ionisation Energy • Why is the 2nd ionisation energy for Na larger than the 1st ionisation energy for the Noble gas in the same period? • Na+→ Na2+ + e 4560 kJ mol-1 • 2)8 2)7 • Ar → Ar+ + e 1530 kJ mol-1 • 2)8)8 2)8)7 • The extra electron level screens (shields) the outer electrons reducing the ionisation energy

33. Questions on Ionisation Energy Answers might also depend on looking at the total number of electrons in the atom Why is there no 4th ionisation energy for Li? It only has 3 electrons!

34. Questions on Ionisation Energy • Each energy value represents one (mole of) electron(s) (need to add the energies if more than one electron removed) • Look at the electron arrangement • Which shell are you removing electrons from (e.g. why is the 2nd IE so high for alkali metals compared to the 1st? IE) • Total number of electrons (e.g. why does Li not have a 4th IE?)

35. Attraction for Bonding Electrons • Different elements have different attractions for bonding electrons • The aim is to obtain a full outer electron shell • Which elements will have the greatest attraction? • Those which require only one or two electrons to fill the outer shell • Those with the smallest atoms – greater “pull” from the nucleus

36. Electronegativity Learn this!! “Measure of the attraction that an atom has for the shared electrons in a covalent bond” • Each element has an electronegativity value from 0 to 4 (the Pauling scale pg 11 of data booklet) • 4 is the highest electronegativity • This element has the greatest attraction for bonded electrons • Plot electronegativity against atomic number for the first 20 elements

37. Start of a new row Electronegativity decreases down a group Electronegativity increases across a period

38. Electronegativity Trends Learn this!! • Going across a period the electronegativity increases, due to • Increased nuclear charge • Decreased atomic size • Greater “pull” on shared electrons • Going down a group the electronegativity decreases, due to • Screening effect of added electron shells • Larger atomic size • Reduced “pull” on shared electrons

39. Increasing Electronegativity The nuclear charge increases and electrons are added to the same energy level. Energy levels are pulled closer to the nucleus. Decreasing Electronegativity The number of energy levels increases. The nuclear charge increases but the outer energy levels are shielded from the effect of the nucleus by the inner energy levels.

40. Electronegativity • The difference in electronegativity values tells us about the type of bonding between atoms Li = 1.0 F = 4.0 F is the most electronegative atom. The difference in electronegativity is 3. This is a large difference and so F removes an electron from Li. The bonding is ionic.

41. Electronegativity • The difference in electronegativity values tells us about the type of bonding between atoms P = 2.2 F = 4.0 F is the most electronegative atom. The difference in electronegativity is 1.8. This is a smaller difference and so F shares an electron P. The pair of electrons will be held closer to the F atom The bonding is (polar) covalent

42. Electronegativity • The difference in electronegativity values tells us about the type of bonding between atoms H = 2.2 P = 2.2 Both atoms have an equal electronegativity. The difference in electronegativity is 0. This mean each atom has an equal pull on the electrons in the bond The bonding is (pure) covalent.

43. Pure Covalent Bonding • Electrons equally shared • No difference in electronegativity • E.g. Cl and Cl • 3.0 and 3.0

44. Ionic Bonding • Electron(s) totally transferred from metal to non-metal • E.g. NaCl • Na = 0.9 • Cl = 3.0

45. Trends can all be found using the data booklet! Learn this!! Decreasing covalent radius Increasing electronegativity Increasing ionisation energy (Increasing nuclear charge + energy levels held closer to the nucleus) Increasing covalent radius Decreasing electronegativity Decreasing ionisation energy (shielding effect of extra energy levels + outer electrons are further from the nucleus)

46. Trends in the Periodic TableStand and Deliver • Work in groups of 4, number yourselves 1-4. • You will each be given an information card about a trend. • The development of the periodic table (easiest) • Atomic size (covalent radii) • Electronegativity • Ionisation energy • To begin, number 1’s should work with another number 1, 2’s with other 2’s and so on. Read over your trend and make a summary in your notes. Compare your notes with the others and make any changes. • Return to your original group and then take it in turns to explain your trend to the group. Number 1’s will go first, then 2 then 3 then 4. • Write a summary in your notes for the other 3 trends. • Complete the card sort to test your knowledge. • Complete the questions.

47. Summarising…… For each trend you should make notes on the following; • A definition of the trend. • The effect down a group of the trend. • The effect across a period of the trend • An explanation of the trends down a group and across a period

49. Answers to questions (Higher Chemistry CfE Chapter 3) • D • B • B • (a) Electronegativity (b) covalent radius decreases across a period from sodium to argon. • (a) the ionic radius for element 13 is less than that of Mg2+ but more than that of B3+. The ionic radius for element 15 is greater than that of N3- (b) (i) H+ has no electrons (ii) Li+ has more protons than H- and therefore a greater nuclear attraction for the same number of electrons. (c) N3- has one more electron energy level. • (a) Chlorine has a greater nuclear charge. (b) Si4+ has 2 layers of electrons, P3- has 3 layers of electrons

50. Chemical Dictionary