The Chemistry of Titanium

1. The Chemistry of Titanium 1b – Trends in the Periodic Table

2. Learning Intentions • To examine trends in properties when moving across or down the periodic table • To be able to define the terms Co-valent Radius, 1st Ionisation Energy, Electronegativity • To be able to explain the trends with regard to atomic structure • To relate nuclear charge, no.of electrons, effect of increasing number of electron shells ( energy levels), size of atom to particular trends

3. 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 on page 5 of the data book

4. 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

5. 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

6. 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.

7. 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.

8. 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.

9. 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 on page 10 of the data book

10. 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

11. 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

12. 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

13. 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

14. Cs Li - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - 1. More Energy Levels 2. Screening Effect As we saw with atomic size, the single electron in the outermost energy level is much further from the nucleus in caesium than in lithium. Caesium’s attraction for its outer electron is lowered by the screening effect caused by all its other electrons.

15. Increasing Ionisation Energy Ionisation Energy Summary Across a period from left to right ionisation energy increases This is due to the increase in atomic charge having a greater pull on the electrons and therefore more energy is required to remove electrons.

16. Increasing Ionisation Energy Decreasing Ionisation Energy Ionisation Energy Summary Down a group ionisation energy decreases This is due to the outer electrons being further away from the nucleus and so the attraction is weaker and they are more easily removed.

17. e e C H Electronegativity Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond Which atom would have a greater attraction for the electrons in this bond and why?

18. Linus Pauling Linus Pauling, an American chemist (and winner of two Nobel prizes!) came up with the concept of electronegativity in 1932 to help explain the nature of chemical bonds. Today we still measure electronegativities of elements using the Pauling scale. Since fluorine is the most electronegative element (has the greatest attraction for the bonding electrons) he assigned it a value and compared all other elements to fluorine. Values for electronegativity can be found on page 10 of the data book

19. Electronegativities Looking across a row or down a group of the periodic table we can see a trend in values. We can explain these trends by applying the same reasoning used for ionisation energies.

20. IncreasingElectronegativity Looking across a period F C B Li Be N O 2.0 2.5 3.0 3.5 4.0 1.0 1.5 What are the electronegativities of these elements? Across a period electronegativity increases The charge in the nucleus increases across a period. Greater number of protons = Greater attraction for bonding electrons

21. F DecreasingElectronegativity Cl Br I Looking down a group 4.0 3.0 What are the electronegativities of these halogens? 2.8 2.6 Down a group electronegativity decreases Atoms have a bigger radius (more electron shells) The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells.

22. 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.