1. The Periodic Law Chapter 5

2. History of the Periodic Table Before 1860, there was no method for accurately determining an element’s atomic mass. Different chemists used different atomic masses for the same elements, resulting in different compositions being proposed for the same compounds.

3. In 1860 scientists led by Stanislao Cannizzaro standardized calculations for atomic mass. Then in 1869, a Russian chemist, Dimitri Mendeleev published the first periodic table. The table was a list of known elements arranged by increasing atomic mass.

4. Mendeleev’s grouped elements showed repeating patterns of properties. Repeating patterns are called periodic functions and the elements arranged in this way are said to show periodicity.

5. The second hand of a watch, for example, passes over and given mark at periodic, 60-second intervals.

6. Mendeleev’s table had blanks which showed places where unknown elements were later placed. Mendeleev also noted discrepancies between grouped properties and some of the element's atomic mass orders.

7. For example: Mendeleev placed iodine (127) after tellurium (128) so tellurium would be in a group of elements with which it shared similar properties. But the question remained—why could most of the elements be arranged in the order of increasing atomic mass, but a few could not?

8. In 1911, English scientist, Henry Mosely (1887-1915) noticed a better pattern. He made a periodic table ordered by increasing atomic number. This modification helped correct the discrepancies between properties and order noted by Mendeleev.

9. Modern Periodic Table

10. Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic number. In other words, when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

11. Periodic Table An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

12. Electron Configuration and the Periodic Table

13. Before we continue… Chemical compounds are formed because electrons are lost, gained, or shared between atoms. The electrons that interact in this manner are those in the highest energy level. These electrons are the most subject to the influence of nearby atoms.

14. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. The stability of the noble gases results from their special electron configurations. The highest occupied energy levels contain stable octets, or 8 valence electrons.

15. Horizontal rows or periods represent primary energy levels. The vertical columns (groups or families) are arranged to place elements with the same outer level electrons (valence electrons) together.

16. Generally the electron configuration of an atom’s highest occupied energy level governs the atom’s chemical properties. So, if the periodic table is arranged to properties, it makes sense that you will see a pattern in the electron configuration.

17. Blocks of the Periodic Table The structure of the table results in blocks of elements based on the filling of the energy sublevels.

18. s-block Elements Include Group 1 (alkali metals) and the Group 2 (alkaline–earth metals). Dissolved in H 2 O to make bases These elements are chemically reactive metals which do not occur as free elements in nature. s-block elements have 1 or 2, valence electrons.

19. Hydrogen Hydrogen is a big exception to the block structure of the periodic table. For convenience hydrogen is usually placed on top of Group 1 although it is not considered an alkali metal.

20. It has a 1s 1 electron configuration but does not share the properties of the elements in group 1. Hydrogen's structure and properties make it unique—it doesn’t really fit with any group.

21. Helium Helium has a similar electron configuration as the Group 2 elements, but it is part of Group 18. Because its highest occupied energy level is filled by 2 electrons, helium possesses special chemical stability like the noble gases.

22. p-Block Elements Elements filling the p sublevel. Includes metals, metalloids and nonmetals. Group 16 (Calcogens) – found in metal ores Includes the most reactive nonmetals, Group 17 (halogens – salt formers). Includes the least reactive elements, Group 18 (noble gases).

23. P-block and s-block elements together make up the main-group or representative elements. All p-block elements have 3-8 valance electrons with the exception of helium, which has 2 valance electrons. The valence electrons are in the outer most s and p sublevels.

24. d-Block Elements All elements in d-block are transition elements. All d-block elements are metals with varying properties caused by interactions of an unfilled d sublevel interacting with an unfilled higher primary energy level. This is where the exceptions to the rules are found. Give gemstones their color

25. f-Block Elements Elements filling the f sublevel. f-block elements should appear in the center of the periodic table, between groups 3 and 4 in the 6 th and 7 th periods. For convenience scientists place these two periods underneath the table to shorten the table.

26. The top row of the f-block is called the lanthanide series. The bottom row of the f-block is called the actinide series. Each series is named after the element that precedes them in the chart (lanthanum- lanthanide) and (actinium- actinide).

27. Trends in the Periodic Table Atomic Radii Ionization Energy Ionic Radii Electron Affinity Electronegativity

28. Atomic Radii Atomic radius is the radius of an atom. It can be defined as one-half the distance between the nuclei of identical atoms that are bonded together. Atomic radii tend to decrease across the period due to an increase in positive nuclear charge. Atomic radii tend to increase going down a group due to addition of a primary energy level.

29. Ionization Energy Ionization energy is the energy required to remove one electron from the neutral atom of an element. An ion is an atom or group of bonded atoms that has a positive or negative charge. Any process that results in the formation of an ion is referred to as ionization.

30. Ionization energy tends to increase across a period. Ionization energy tends to decrease down a group. The energy to remove the first electron is called the first ionization energy. Removing successive electrons requires increased energy.

31. Electron Affinity Neutral atoms can also acquire electrons, not just lose them. The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.

32. Most atoms release energy when they acquire electrons. Energy released is represented as a negative number, energy absorbed by a positive number. Think of energy as a balance in a bank account…

33. As you move across a period, it is generally easier for the atoms to acquire electrons. The values become increasingly negative because the easier it is for the atom to gain an electron, the more energy will be released.

34. As a general rule, electrons add with greater difficulty going down a group, but there are exceptions. The main thing you need to know is what electron affinity is.

35. Ionic Radii Ionic radius is the radius of an ion. Includes positive charged atoms (cations) and negatively charged atoms (anions). Ionic radii tend to decrease across a period due to increase in positive charge to negative charge ratio. Ionic radii tend to increase down a group due to the addition of a primary energy level.

36. Electronegativity In many compounds, one element will attract the electrons more strongly than the other. The uneven concentration of charge has a significant effect on the chemical properties of a compound and therefore it is useful to have a measure of how strongly one atom attracts the electrons of another atom within a compound.

37. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons. Electronegativity tends to increase across a period. Electronegativity tends to decrease (or remain the same) down a group. The element with the highest electronegativity is F and the lowest electronegativity is Fr.