Topic 3: Periodicity

1. Topic 3: Periodicity

2. Topic 3: Periodicity 3.1 The periodic table 3.1.1      Describe the arrangement of elements in the periodic table in order of increasing atomic number 3.1.2      Distinguish between the terms group and period 3.1.3      Apply the relationship between the electron arrangement of elements and their position in the periodic table up to z=20. 3.1.4      Apply the relationship between the highest occupied energy level for an element and its position in the periodic table.

3. 3.1.1.Describe the arrangement of elements in the periodic table in order of increasing atomic number. Development of the Periodic Table • Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

5. Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913) Elements arranged by increasing atomic number into • periods (rows) and • groups or families (columns), which share similar characteristics

6. Groups or Families: Vertical Lines • Elements in the same group have similar chemical and physical properties!! • Why? • They have the same number of valence electrons. • They will form same kind of ions. • Combine the same way

7. Family Names: • Group 1: alkali metals • Group 2: alkaline earth metals • Transition metals • Group 7: halogens • Group 8/0: noble gases

9. Periods: horizontal row (7). Rows in the periodic table are called periods. • As one moves from left to right in a given period, the chemical properties of the elements change.

10. Elements in the same row or period, have same number of energy levels. 11Na 13 Al 15P

12. Periodic Table & Electronic Configuration

13. ns2np6 ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f Ground State Electron Configurations of the Elements

14. Lewis Dot Diagram Symbol of the element and dots representing the valence electrons Na ● Ca ● ● Al ● ● ●

17. Shielding • Shielded slightly from the pull of the nucleus by the electrons that are in the closer orbitals.(inner e)

18. Effective Nuclear Charge • Effective nuclear charge is the charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.(Effective nuclear charge is essentially the positive charge that a valence electron "sees“) • http://dl.clackamas.cc.or.us/ch104-06/efffective_nuclear_charge.htm

20. 3.2 Physical properties3.2.1      Define the terms first ionization energy and electronegativity3.2.2      Describe and explain the trends in atomic radii, ionic radii, first ionization energy, electronegativities and melting points for alkali metals (Li  Cs) and the halogens (F  I).3.2.3      Describe and explain the trends in atomic radii, ionic radii, first ionization energy, and electronegativities for elements across period3.2.4      Compare the relative electronegative values of two or more elements based on their position on the periodic table.

21. Video • http://www.youtube.com/watch?v=-4xKhr8RNjA

22. Atomic Size, Radii • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time( diatomic) • Summary: it is the volume that an atom takes up http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf

23. Group trends H • As we go down a group (each atom has another energy level) Li Na K Rb

24. Periodic Trends The atomic radius decreases as you go from left to right across a period. • Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter. Remember filling up same energy level, little shielding occurring. Na Mg Al Si P S Cl Ar

25. Ionic Size - Cations Cations form by losing electrons. Cations are smaller than the atom they come from. . Ca+2 < Ca

26. Ionic size - Anions • Anions form by gaining electrons. • Anions are bigger than the atom they come from. N-3 > N

27. Periodic Trends • Metals losing from outer energy level, more protons than electrons so more pull, causing it to be a smaller species. • Non metals gaining electrons in its outer energy level, but there are less protons than electrons in the nucleus, so there is less pull on the protons, so found further out making it larger. N-3 B+3 O-2 F-1 Li+1 C+4 Be+2

28. Size of Isoelectronic ions • Positive ions have more protons so they are smaller. N-3 O-2 F-1 Ne Na+1 Al+3 Mg+2

29. Electronegativity

30. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair it shares. • Big electronegativity means it pulls the electron toward it. • Atoms with large negative electron affinity have larger electronegativity.

32. Group Trend • The further down a group the farther the electron is away and the more electrons an atom has. • So as you go from fluorine to chlorine to bromine and so on down the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

33. Period Trend • Electronegativity increases from left to right across a period • When the nuclear charge increases, so will the attraction that the atom has for electrons in its outermost energy level and that means the electronegativity will increase

34. Period trend Electronegativity increases as you go from left to right across a period. • Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

35. Group Trend electronegativity decreases as you go down a group. • Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. • This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

36. Melting Points of Group 1

37. Metallic bonding • Collective bond, not a single bond • Strong force of electromagnetic attraction between delocalized electrons (move freely). • This is sometimes described as "an array of positive ions in a sea of electrons

38. Why does the melting point decrease going down the alkali metals family? • Atoms are larger and their outer electrons are held farther away from the positive nucleus. • The force of attraction between the metal ions and the sea of electrons thus gets weaker down the group. • Melting points decrease as less heat energy is needed to overcome this weakening force of attraction.

39. Melting Points for halogens

41. Why does melting point increase going down the halogens? • The halogens are diatomic molecules, so F2, Cl2, Br2, I2 • As the molecules get bigger there are more electrons that can cause more influential intermolecular attractions between molecules. • The stronger the I.A, the more difficult it will be to melt. (more energy needed to break the I.A)

42. What are these I.A? van der Waals forces (London dispersion): • Electrons are mobile, and although in a diatomic molecule they should be shared equally, it is found that they temporarily move and form slightly positive end and negative end. • Now that one end is + and the other -, there can be intermolecular attractions between the opposite charges of the molecules

43. van der Waals forces

44. IB requires knowledge specifically for halogens. Check out this site for more detail. http://www.chemguide.co.uk/inorganic/group7/properties.html

45. Period 3 melting point trends

46. Explanation • M.P rise across the 3 metals because of the increasing strength of the metallic bonds. • Silicon has a giant covalent structure just like diamond which makes its structure remarkably strong and therefore takes more energy to break apart.

47. The atoms in each of these molecules are held together by covalent bonds (except Ar) • They would have weak I.A affecting the amount of energy needed to melt them. • Ar has extremely weak forces of attraction between its atoms, so its easiest to melt.

48. 3.3 Chemical properties 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. 3.3.2 Discuss the changes in nature from ionic to covalent and from basic to acidic of the oxides across period 3

49. Reactivityof alkali metals • Generally group 1 metals become more reactive as you go down a group. • The valence electron of group 1 are found further from the nucleus as you go down the group. • It is easier to remove an electron from francium than from lithium

50. Alkali metal + water • Li(s) + H2O (l)  LiOH(aq) + H2 (g) (Li + and OH- in solution) • The metal reacts with water to form the hydroxide of the metal (strong base) and bubbles off hydrogen gas. • The larger the alkali metal, the more vigorous the reaction. Sometimes the H2 gas actually lights itself (exothermic reaction, releases heat) causing the H2 to burn.