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  1. Unit 2 – 2014-2015 Periodicity

  2. Lesson 1: Review of Periodic Table Structure Thursday, October 2nd

  3. Periodicity • IB Understandings • The periodic table is arranged into four blocks associated with the four sublevels – s, p, d and f • The periodic table consists of groups (vertical columns) and periods (horizontal columns) • The period number (n) is the outer energy level that is occupied by electrons • The number of the principal energy level and the number of the valence electrons in an atom can be deduced from its position on the periodic table • The periodic table shows the positions of metals, non-metals and metalloids

  4. 3.1 Applications and Skills • Applications and skills: • Deduction of the electron configuration of an atom from the element’s position on the periodic table, and vice versa. • Guidance: • The terms alkali metals, halogens, noble gases, transition metals, lanthanoids and actinoids should be known. • The group numbering scheme from group 1 to group 18, as recommended by IUPAC, should be used.

  5. The development of the periodic table brought a system of order to what was otherwise an collection of thousands of pieces of information. The periodic table is a milestone in the development of modern chemistry. It not only brought order to the elements but it also enabled scientists to predict the existenceof elements that had not yet been discovered . Periodic Table

  6. Early Attempts to Classify Elements • Dobreiner’sTriads (1827) • Classified elements in sets of three having similar properties. • Found that the properties of the middle element were approximately an average of the other two elements in the triad. 6

  7. Dobreiner’sTriads Note: In each case, the numerical values for the atomic mass and density of the middle element are close to the averages of the other two elements

  8. Newland’s Octaves -1863 • John Newland attempted to classify the then 62 known elements of his day. • He observed that when classified according to atomic mass, similar properties appeared to repeat for about every eighth element • His attempt to correlate the properties of elements with musical scales subjected him to ridicule. • In the end his work was acknowledged and he was vindicated with the award of the Davy Medal in 1887 for his work. .8

  9. Dmitri Mendeleev Dmitri Mendeleev is credited with creating the modern periodic table of the elements. He gets the credit because he not only arranged the atoms, but he also made predictions based on his arrangements His predictions were later shown to be quite accurate. .9

  10. Mendeleev’s Periodic Table • Mendeleev organized all of the elements into one comprehensive table. • Elements were arranged in order of increasing mass. • Elements with similar properties were placed in the same row. .10

  11. Mendeleev’s Periodic Table

  12. Mendeleev’s Periodic Table Mendeleev left some blank spaces in his periodic table. At the time the elements gallium and germanium were not known. He predicted their discovery and estimated their properties.

  13. Periodic Properties • Elements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC Periodic properties include: -- Ionization Energy --Electronegativity -- Electron Affinity -- Atomic Radius -- Ionic Radius .13

  14. REVIEW: Periodic Table • Elements are arranged by increasing atomic number, Z • Groups, or vertical columns, group elements with the same number of valence electrons and therefore similar chemical properties • Periods, or rows, are horizontal groups; the period number is equal to the principal quantum number, n, of the highest occupied energy level of the elements in that period

  15. Important Group To Remember

  16. Arrangement • Metals make up most of the periodic table and are found on the left side • Metalloids include B, Si, Ge, As, Sb and Te and separate metals and non-metals (Po and At are sometimes considered metalloids); have characteristics of both metals and nonmetals • Non-Metals are on the right side of the periodic table after the metalloids

  17. Metals and Nonmetals Metalloids Transition metals NONMETALS METALS METALS

  18. Additional Groupings in the Periodic Table • Nonmetals, Metals, Metalloids, Noble gases

  19. IB Goodness! • You have a large number of periodic tables in your data booklet • All groups are numbered 1-18 • The position of an element is related to its electron configuration • Know the s-block, p-block, d-block and f-block; those blocks represent which valence electrons are getting filled

  20. Blocks p d S f

  21. Let’s Practice!

  22. Let’s Practice

  23. Lesson #2 – Review of Periodic Table Families Friday, October 3rd

  24. 3.2 Periodic Trends IB Understandings • Vertical and horizontal trends in the Periodic Table exist for atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity. Guidance • Only examples of general trends across periods and down groups are required. For ionization energy the discontinuities in the increase across a period should be covered. • Trends in metallic and non-metallic behaviour are due to the trends above. • Oxides change from basic through amphoteric to acidic across a period.

  25. Topic 3.2 Applications and Skills Applications and skills: • Prediction and explanation of the metallic and non-metallic behaviour of an element based on its position in the Periodic Table. • Discussion of the similarities and differences in the properties of elements in the same group, with reference to alkali metals (Group 1) and halogens (Group 17). Guidance • Group trends should include the treatment of the reactions of alkali metals with water, alkali metals with halogens and halogens with halide ions. • Construction of equations to explain the pH changes for reactions of Na2O, MgO, P4O10, and the oxides of nitrogen and sulfur with water.

  26. Periodicity • Properties of elements repeat themselves periodically • The Periodic Table is arranged to show these trends

  27. Effective Nuclear Charge • Nuclear charge is the number of protons in the nucleus of an atom and increases steadily from left to right on the Periodic Table • HOWEVER, the outer electrons are shielded from the nuclear charge by the inner electrons and feel less of this positive pull • The effective nuclear charge experience by the outer electrons is less than the full nuclear charge

  28. Effective Nuclear Charge - Trends • Effective nuclear charge INCREASES from left to right across the Periodic Table (no change in number of inner electrons) • Effective nuclear charge REMAINS THE SAMEas we go down a group

  29. Atomic Radius • Atomic radiusis measured as ½ the distance between neighboring nuclei • Atomic radii DECREASEfrom left to right across the period and INCREASE down a group

  30. Atomic Radius (cont.) • Atomic radii increase down a group because we are adding Principle Energy Levels • Atomic radii decrease from left to right as we increase effective nuclear charge

  31. Ionic Radius • Positive ions (cations)have a smaller radius than the parent atom; lose outer valence shell • Negative ions (anions) have a larger radius than the parent atom; adding electrons to outer shell which increases electron repulsion • Atomic radii decrease from Groups 1 to Group 14 for the positive ions as we increase effective nuclear charge • Atomic radii decrease from Groups 14 to Group 17 for the negative ions as we increase effective nuclear charge • Ionic radii increase down the group as number of energy shells increases

  32. Let’s Practice • Describe and explain the trend in radii of the following atoms and ions: O2–, F–, Ne, Na+, and Mg2+. • The ions and the Ne atom have 10 electrons and the electron configuration 1s22s22p6. The nuclear charges increase with atomic number: • O: Z = +8, F: Z = +9, Ne: Z = +10, Na: Z = +11, Mg: Z = +12 • The increase in nuclear charge results in increased attraction between the nucleus and the outer electrons. The ionic radii decrease as the atomic number increases.

  33. Ionization Energies

  34. Ionization Energy Trends • Ionization energy increases across a period as we increase effective nuclear charge • Ionization energydecreases down a group as we increase the shielding effect • Exceptions to these trends can be explained in terms of electron configuration stability (i.e. it takes more energy to remove an electron from a full sub-level or shell)

  35. Electron Affinity • The first electron affinity of an element is the energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions: X(g) + e– → X–(g) • Values are tabulated in Table 8 of the IB data booklet. • 1st electron affinity is general exothermic as the electron is attracted to the positive nucleus; 2nd and 3rd electron affinities become endothermic as the atom already has a negative charge!

  36. Electron Affinity Trends • Group 17 have the highest electron affinity • Group 1 has the lowest electron affinity • Group 2 and Group 15 have the highest electron affinities; here you are adding the first electron to a half-filled sub-level; electrostatic repulsion

  37. Electronegativity • Theelectronegativityof an element is a measure of the ability of its atoms to attract electrons in a covalent bond • Electronegativity increases from left to right across a period owing to the increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons. • Electronegativity decreases down a group. The bonding electrons are furthest from the nucleus and so there is reduced attraction.

  38. Metals vs. Non-Metals • The ability of metals to easily conduct electricity is due to their low ionization energy and ability to move their electrons away from the nucleus • As you move from left to right on the Periodic Table, there is a slow transition from metal to semi-metal to non-metal

  39. Lesson #3 – Review of Periodic Table Families Tuesday, October 7, 2014

  40. Nota Bene • Remember when studying that the group of an element on the Periodic Table also tells you the number of valence electrons for that element!

  41. The Electron Shielding Effect • Electrons between the nucleus and the valence electrons repel each other making the atom larger. .41

  42. How is atomic radius measured? • Half the distance between neighboring nuclei. • OR • Distance from the nucleus to the outermost electron.

  43. Atomic Radius Across a Period • Why does atomic radius decrease across a period? • Think about the effective nuclear charge! How many occupied energy levels does each atom have? Fluorine’s radius is almost half of Lithium’s.

  44. Atomic Radius

  45. Atomic Radius • Extra Info: • We typically measure atomic radius at bonding atomic radius where we look at atoms in chemical bonds with other atoms of the same element and take ½ the diameter between the two nuclei • There is also nonbonding atomic radius or van der Waals’ radius for things like Noble Gases that do not bond; take a look at these atoms in the solid phase and measure the distance between two nuclei

  46. Ionic Radius – Some More Info • The factors that effect ionic radius include nuclear charge, number of filled energy shells, electrostatic repulsion and overall charge • At the HL level, the IB likes to ask questions about ionic radius that are not clear-cut and require you to think • Which ion would have a larger ionic radius, Na+ or Mg2+? Let’s think about this…we cannot directly compare these thinking about trends from left to right because they have different charges! But, we can recognize that both now have the electron configuration of Neon. However, Mg has one more proton in the nucleus meaning electrons are pulled in tighter!

  47. Atomic Radius .47

  48. Trends in Ion Sizes Radius in pm .48

  49. Effective Nuclear Charge • Extra Info (way beyond IB!) • The formula for effective nuclear charge is: Zef = Z – S Where s= screening or shielding constant

  50. Melting Points • Comparing melting points is complex because it depends on bonding as well as nuclear charge • Trends down Groups 1 through 17 can be explained as elements in each group bond in the same way! • Melting pointsDECREASEdown Group 1 as these metallic bonds have delocalized electrons and as you move down the group the attraction decreases • Melting points INCREASE down Group 17 because these diatomic elements are held together by London Dispersion forces which get stronger as the electron cloud gets bigger • Melting points generally rise from left to right until Group 14 and then fall from Group 14 to Group 18