Trends in the Periodic Table and Bonding

1. Using your data book and using your knowledge of chemistry comment on the use of titanium metal in the aerospace industry.

2. Trends in the Periodic Table and Bonding

3. I can predict the properties of an element using its position in the Periodic Table I can identify groups and periods in the Periodic Table. I can explain why certain elements have similar properties I can identify the alkali metals, halogens, noble gases and transition elements in the Periodic Table. Arrangement of Elements in the Periodic Table

4. Data Book Task • Use your iPad to join the Higher Chemistry class on iTunesU. Enrolment key for the iTunesU course is: FKH-NPF-BXL • Download a copy of the Higher Chemistry Data book either from the SQA website or the iTunesU resource folder. • Using your data book and using your knowledge of chemistry comment on the use of titanium metal in the aerospace industry.

5. The Periodic Table • On March 6th 1869 Dmitri Mendeleev, a Russian chemist, published his Periodic Table of the Elements. He arranged the known elements in order of increasing atomic masses. This was eventually changed to atomic number. • Elements with similar chemical properties were arranged in groups. He left gaps for elements yet to be discovered • In the years that followed his ideas were modified and new elements were discovered until we arrived at the modern Periodic Table.

6. Trends in the Periodic Table The elements in the periodic table have different properties The table is set up in such a way that these properties vary periodically across a period, or down a group • The properties are both physical and chemical • The chemical properties of an element stem from its physical properties: • Density, melting points and boiling points, atomic size, ionisation enthalpy, attraction for bonding electrons

7. After discussion with your group, make sure you can identify where all the following groupings are in the Periodic Table and what their properties are. Mark these on the periodic table handout and stick it into your notebook. Metals Non-metals Alkali metals Transition metals Halogens Noble gases The diatomic elements The radioactive elements

8. The diatomic elements Noble gases Alkali metals Halogens Transition metals Metals Non-metals

9. Physical properties of the elements There are variations in the physical properties of the elements across a period and down a group. (i) Density • Copy and complete the table below for the first twenty elements, using the information in the data booklet. • What do all the elements in Group A have in common? • What do all the elements in Group B have in common? • How does the density of the elements change across a period? • How does the density of the elements change down a group?

10. Variation of density (g cm-3) with atomic number period 2 (Li - Ne) maximum at boron (B) - group3 period 3 (Na - Ar) maximum at Aluminium (Al)- group 3 Al B Na Li Ne Ar

11. Variation of density (gcm-3) with atomic number In general in any period of the table, density first increases from group 1 to a maximum in the centre of the period, and then decreases again towards group 0 5th 4th 3rd 2nd

12. Variation of density (g cm-3) with atomic number Adapted from New Higher Chemistry E Allan J Harris down a group gives an overall increase in density In Ga Al B Cs Rb Na K Li

13. Melting and Boiling points The melting and boiling points of elements give an indication of the forces that hold the atoms or molecules together The higher the melting and boiling point the stronger the forces The trend is similar for both melting and boiling so we’ll just look at melting

14. (ii) Melting and boiling points • Using the information in the data booklet, how do the melting and boiling points of the elements change across a period? • What is the trend in melting and boiling points down Group 1? • What is the trend in melting and boiling points down Group 7 and Group 0? • The melting point starts off low, gradually increases to a peak (at group 4) then gradually decreases to a very low value (at group 0 or 8) • To explain this trend we must think about the strength of the forces between the molecules • In group 1 the atoms are held together by metallic bonds • In group 4 the atoms are held together by many very strong covalent bonds (covalent network) • In group 8 the atoms are held together by very weak bonds (monatomic gases) • We will look at the different types of bonding later in the unit

15. I can explain how a covalent bond is formed. I can describe the behaviour of outer electrons in metallic bonding. I can explain the difference between covalent network and covalent molecular. I can give examples of metallic, covalent molecular, covalent network and monatomic elements. Bonding and Structure of the first twenty elements

16. N5 2014 PP Question

17. Bonding in the first 20 elements

18. Demo 1.6

19. 1. Metallic Bonding e.g. Li, Na, K, Be, Mg, Ca, Al Strong electrostatic forces exist between the positive nuclei and the outer shell electrons. These electrostatic attractions are known as metallic bonds. + + + + Positive nucleus (core) Delocalised electron + + + + • The outer shell in metals is not full and so metal electrons can move between these partially filled outer shells. • This creates what is sometimes called a ‘sea’ or ‘cloud’ of delocalised electrons.

20. + Physical properties of metals Metals are malleable and ductile applied force Metal atoms can ‘slip’ past each other because the metallic bond is not fixed and it acts in all directions. 2. Conduction of electricity The ‘sea’ of delocalised electrons can move and carry the charge 3. Change of state M.p.’s are relatively low compared to the B.P’s. This is because in a molten metal the metallic bonding is still present. B.p.’s are much higher as you need to break the metallic bonds throughout the metal lattice. • Metal b.p.’s are dependant on • How many electrons are in the outer shell • How many electron shells there are.

21. 2. Covalent Networks Giant lattice of covalently bonded atoms. e.g. B, C, Si Giant molecules held together by covalent bonds, resulting in high mpt and bpt. Boron: M.pt.= 2573K= 2300oC Silicon: M.pt.= 1683K= 1410oC Model 1.8

22. Two Forms of Carbon 1. Graphite (covalent network) Strong covalent bonds between atoms Weak Van der Waals forces between layers

23. 3 electrons from each C atom used in bonding. 1 electron from each C is delocalised so graphite conducts electricity. Stacked hexagonal layers of C atoms with only weak Van der Waals between layers so layers can slip and slide- graphite is soft and can be used as a lubricant.

24. 2. Diamond (covalent network) Hardest natural substance  very strong covalent bonds. Electrical insulator (no delocalised electrons)-all outer electrons are localised in covalent bonds. Tetrahedrally bonded carbon atoms We will discuss the 3rd form of carbon later

25. 3. Discrete Covalent Molecules These molecules have known numbers of atoms (discrete molecules). e.g. Oxygen M.pt.= 55K= -218oC Sulphur M.pt.= 386K= 113oC Low Mpts indicate that weak Van der Waals forces are present.

26. } Diatomic a) Diatomic molecules H – H O = O N ≡ N F – F Cl – Cl Remember HON 7!  All gases due to weak Van der Waals forces.

27. Covalent solids held together by Van der Waals which are stronger due to higher molecular masses. b) P and S

28. c) Carbon Buckminster Fullerene Very large, C60.

29. 4. Monatomic Group 8 elements e.g. He, Ne, Ar Non-bonded atoms. Only weak Van der Waals forces between atoms. Noble gases have full outer shells, they do not need to combine with other atoms.

30. -100 -120 -140 -160 -180 -200 -220 -260 -280 Noble gases Helium Neon Argon b.p / oC Krypton Xeon B.p.’s increase as the size of the atom increases This happens because the Van der Waals forces increase

31. PP Question 2004

33. Bond Strengths

34. Bonding and melting point Expt 1.9 The properties of a compound depend on the type of bonding present.  In the following experiment you will investigate melting points of ionic and covalent molecular solids. Place the test tubes provided in a beaker of boiling water for a few minutes Find out the actual melting point of these compounds.

35. Ionic compounds have ……………. melting points. • Explain this in terms of arrangement and movement of particles as well as attraction between particles. Covalent molecular solids have …………………… melting points. • Explain this in terms of attraction between and movement of particles.

36. Covalent network compounds Silicon dioxide has the formula SiO2. Silicon carbide has the formula SiC. • What type of bonding would you expect to exist in these compounds? • Would you expect these compounds to have high or low melting points? • Find out the actual melting points. • How do they fit with your prediction? • Examine models of these compounds. • Explain melting point and boiling point in terms of bonding and movement of particles • Consider the valances of C and Si and use this information to work out the exact structure of silicon carbide. • Silicon carbide (SiC) is widely used as an abrasive as it is an extremely hard material. Explain this is terms of its structure.

37. I can explain how a covalent bond is formed. I can describe the behaviour of outer electrons in metallic bonding. I can explain the difference between covalent network and covalent molecular. I can give examples of metallic, covalent molecular, covalent network and monatomic elements. Patterns in the periodic table: Covalent Radius

38. I can use covalent radius to describe the changes in the size of atoms across a period and down a group. I can explain the change in covalent radius in terms of changes in the number of occupied shells or the nuclear charge. I can state what is meant by first, second and third ionisation energies. I can write state equations to represent first, second and third ionisation energies. I can use atomic size and screening effect to explain the change in ionisation energies down a group. I can use atomic size and nuclear charge to explain the change in ionisation energies across a period. Periodic Trends in Ionisation Energy and Covalent radius

39. Covalent Radii of Elements The size of an atom is measured by it’s covalent radius, the distance between the nucleus and it’s outer electrons. nucleus covalent radius energy levels Values for covalent radii can be found in the data book

40. Cs Li - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - Looking down a group The single electron in the outermost energy level is much further from the nucleus in caesium. The caesium atom also has many more electrons between the single outer electron and the nucleus. This causes the caesium atom to have a much larger covalent radius. This screening effect counteracts the attraction from the greater nuclear charge.

41. 3+ 9+ - - - - - - - - Looking across a period Across a period we can see the covalent radius decreasing. As we move left to right we are adding a proton to the nucleus and an electron to the outermost energy level. So, from lithium to fluorine: Lithium Atom Fluorine Atom

42. 3+ 9+ 9+ - - - - - - - - - - - - - - - Looking across a period The lithium atom has a smaller nuclear charge than neon and so a larger covalent radius Fluorine’s greater nuclear charge pulls the outer energy level in closer. radius = 134pm radius = 71pm

43. Decreasing Atomic Size Atomic Size Summary Across a period from left to right atomic size decreases This is because of the atom having more electrons & protons and therefore a greater attraction which pulls the atom closer together hence the smaller size.

44. Decreasing Atomic Size Increasing Atomic Size Atomic Size Summary Down a group atomic size increases This is because of the extra outer energy levels and the screening effect of the outer electrons.

45. Ionisation Energy The ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state. The first ionisation energy of magnesium: Mg (g) Mg+ (g) + e- 744 kJmol-1 Values for ionisation energies can be found in the data book

46. Ionisation Energy The second ionisation energy of magnesium: Mg+ (g) Mg2+ (g) + e- 1460 kJmol-1 The third ionisation enthalpy shows a massive increase because it requires an electron to be removed from magnesium’s second energy level. Mg2+ (g) Mg3+ (g) + e- 7750 kJmol-1

47. Looking across a period From lithium to neon the first ionisation energy increases. Why? B Ne Li Be C N O F Li (g) Li+ (g) + e- 526 kJmol-1 Ne (g) Ne+ (g) + e- 2090 kJmol-1

48. 3+ - An atom of Lithium The lithium atom has 3 protons inside the nucleus The outer electron is attracted by a relatively small nuclear charge Li (g) Li+ (g) + e- 526 kJmol-1

49. 10+ - - - - - - - - An atom of Neon The neon atom has 10 protons inside the nucleus Each of neon’s eight outer electrons is attracted by a stronger nuclear charge Ne (g) Ne+ (g) + e- 2090 kJmol-1

50. Looking down a group The first ionisation energy decreases down a group in the periodic table. Why? Li (g) Li+ (g) + e- 526 kJmol-1 Cs (g) Cs+ (g) + e- 382 kJmol-1